4.04 case study acid rain

• Open access
• Published: 12 October 2021

Effects of sulfuric, nitric, and mixed acid rain on the decomposition of fine root litter in Southern China

• Xin Liu 1 ,
• Miaojing Meng 1 ,
• Yong Zhang 2 ,
• Chong Li 1 ,
• Shilin Ma 1 ,
• Qinyu Li 1 ,
• Qiong Ren 3 ,
• Yinlong Zhang 4 &
• Jinchi Zhang 1  

Ecological Processes volume  10 , Article number:  65 ( 2021 ) Cite this article

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Metrics details

China has been increasingly subject to significant acid rain, which has negative impacts on forest ecosystems. Recently, the concentrations of NO 3 − in acid rain have increased in conjunction with the rapid rise of nitrogen deposition, which makes it difficult to precisely quantify the impacts of acid rain on forest ecosystems.

For this study, mesocosm experiments employed a random block design, comprised of ten treatments involving 120 discrete plots (0.6 m × 2.0 m). The decomposition of fine roots and dynamics of nutrient loss were evaluated under the stress of three acid rain analogues (e.g., sulfuric (SO 4 2− /NO 3 − 5:1), nitric (1:5), and mixed (1:1)). Furthermore, the influences of soil properties (e.g., soil pH, soil total carbon, nitrogen, C/N ratio, available phosphorus, available potassium, and enzyme activity) on the decomposition of fine roots were analyzed.

The soil pH and decomposition rate of fine root litter decreased when exposed to simulated acid rain with lower pH levels and higher NO 3 − concentrations. The activities of soil enzymes were significantly reduced when subjected to acid rain with higher acidity. The activities of soil urease were more sensitive to the effects of the SO 4 2− /NO 3 − (S/N) ratio of acid rain than other soil enzyme activities over four decomposition time periods. Furthermore, the acid rain pH significantly influenced the total carbon (TC) of fine roots during decomposition. However, the S/N ratio of acid rain had significant impacts on the total nitrogen (TN). In addition, the pH and S/N ratio of the acid rain had greater impacts on the metal elements (K, Ca, and Al) of fine roots than did TC, TN, and total phosphorus. Structural equation modeling results revealed that the acid rain pH had a stronger indirect impact (0.757) on the decomposition rate of fine roots (via altered soil pH and enzyme activities) than direct effects. However, the indirect effects of the acid rain S/N ratio (0.265) on the fine root decomposition rate through changes in soil urease activities and the content of litter elements were lower than the pH of acid rain.

Conclusions

Our results suggested that the acid rain S/N ratio exacerbates the inhibitory effects of acid rain pH on the decomposition of fine root litter.

Introduction

Acid rain has emerged as one of the most important global scale environmental challenges (Li et al. 2019 ), and China has become the third most seriously affected regions worldwide to its effects (Liang et al. 2016 ). Furthermore, the composition of acid rain in China has progressively transitioned from sulfuric acid rain (SAR) to nitric acid rain (NAR) (Liu et al. 2020 ), due to the recent regulation of sulfur dioxide (SO 2 ) emissions, and relatively stable nitrogen oxide (NO x ) emissions (Liu and Zhang, 2019 ; Yu and Duan 2020 ). Previous studies have found that the ratios of SO 4 2− /NO 3 − in the precipitation of Beijing, Lin’an, and Guangzhou in China decreased from 4.28 ~ 5.40 to 1.50 ~ 1.81 over the last few decades (Fang et al. 2011 ; Li et al. 2010 ; Wang et al. 2012 ). Currently, although the overall air quality has been mitigated slightly over the last few years as the result of reduced SO 2 and NO x emissions, acid rain remains a challenging issue for China (Liu et al. 2019 ).

Acid rain increases risks to the integrity of forest ecosystem, particularly in terms of its changing compositions (Liu et al. 2018 ). Plant litter, as one of its main components, plays critical roles in the carbon (C) balance and nutrient cycling of forest ecosystems (Penner and Frank 2019 ). Previous studies revealed that the SO 4 2− /NO 3 − ratio of acid rain is a significant factor that profoundly influences litter decomposition processes (Lv et al. 2014 ; Liu et al. 2017a ). However, these earlier investigations set their primary focus on the decomposition of litters on the ground surface (Wang et al. 2010 ; Tang et al. 2019 ) while essentially ignoring the dynamics of fine root decomposition. In contrast to seasonal deciduous leaves, the decomposition of fine roots can occur year-round, which has the function of continuously supplying nutrients to the soil (Xia et al. 2018 ; See et al. 2019 ).

The decomposition of fine root litter is mediated by the litter quality (Couteaux et al. 1995 ; Penner and Frank 2019 ), soil environment (DeForest 2019 ; Tresch et al. 2019 ), and climate (Penner and Frank 2019 ; Allison et al. 2018 ). Litter quality is generally considered to be the dominant control of decomposition (Bradford et al. 2016 ; Chao et al. 2019 ), where plant litters with higher nitrogen (N) and N:C ratios decompose more rapidly (Patoine et al. 2017 ). Acid rain can alter the soil environment, encompassing soil enzymes and the availability of soil nitrogen (Wang et al. 2010 ; Liu et al. 2017a ). These changes modify the loss dynamics of elements and compounds in fine roots (Wang et al. 2010 ), which leads to gradual changes in substrate quality, which subsequently impacts the decomposition process (Liu et al. 2017b ). However, there are few studies that explore the dynamics of element loss in fine roots under changing acid rain compositions.

Soil enzymes play pivotal roles in litter decomposition and nutrient cycling in forest ecosystems (Lv et al. 2014 ), where their activities are the direct expression of soil microbial communities to metabolic requirements and available nutrients (Liu et al. 2020 ; Ling et al. 2010 ). Previous studies revealed that the inhibitory effects of nitric acid rain (NAR) on most enzyme activities were more obvious than those of sulfuric acid rain (SAR) (Lv et al. 2014 ; Liu et al. 2020 ). Nevertheless, there remains a scarcity of data as relates to the impacts of soil enzyme activities on the decomposition of fine root litter under the stress of changing types of acid rain, which seriously affects comprehensiveness and objectivity in the assessment and prediction of global C cycles (Liu et al. 2018 ; Huang et al. 2019 ).

To explore the impacts of changing types of acid rain on the decomposition of fine root litter and element loss dynamics, we established a simulated experiment at a Cunninghamia lanceolata (evergreen coniferous) plantation in China, which comprises ~ 25% of plantations in the subtropical areas of Southern China (Liu et al. 2018 ). Based on previous research (Liu et al. 2017a , 2018 ) we hypothesized that: (1) acid rain inhibits the decomposition of fine root litter and activities of soil enzymes. (2) The negative impacts of simulated acid rain on the decomposition of fine root litter and soil enzyme activities intensify as NO 3 − concentrations increase.

This study was conducted in the Yangtze River Delta Region, which is one of the more seriously affected regions in China in terms of acid rain. The experimental plots were located in a Cunninghamia lanceolata plantation at the Tong Shan Forestry Farm (31°37′N, 118°51′E) in Nanjing, China. The average annual pH of the precipitation in China in 2020 was ~ 5.65, according to the Environment Bulletin of China. Southern China remains a seriously impacted region for acid rain, with the lowest average precipitation pH value reaching 4.39 (Additional file 1 : Fig. S1). The soil type of the C. lanceolata plantation was yellow brown soil and clay with a thickness of ~ 60 cm. The parent material layer is weathered sandstone, whereas the soil moisture content and temperature were 9.5–27.9% and 6.1–28.9 ℃, respectively. The soil pH, total carbon (TC), total nitrogen (TN), total sulfur (TS), available phosphorus (AP), and available potassium (AK) for the 0–10 cm soil depth soil were 4.23 ± 0.12, 38.42 ± 7.91 mg g −1 , 3.55 ± 0.63 mg g −1 , 0.85 ± 0.12 mg g −1 , 2.58 ± 0.68 mg kg −1 , and 35.44 ± 6.25 mg kg −1 , respectively.

Simulated acid rain treatments

A total of 120 discrete plots (0.6 m × 2.0 m) were established in the C. lanceolata plantation, which were separated from each other by ~ 5 m as described by Liu et al. ( 2018 ). Each plot was 1.0 m from the single C. lanceolata , which was above the plot (Additional file 1 : Fig. S2). Ten experimental treatments were formulated, including three categories of acid rain (sulfate acid rain, nitric acid rain, and mixed acid rain), three acid rain acidities (pH 4.5, pH 3.5, and pH 2.5), and a control (CK) treatment (local stream water—pH 6.6). Three replicate plots were randomly selected to receive the simulated acid rain (AR) treatments.

Stock solutions of sulfate acid rain (S), mixed acid rain (SN), and nitric acid rain (N) were prepared by combining 0.5 mol L −1 H 2 SO 4 and 0.5 mol L −1 HNO 3 at (S/N) molar ratios of 5:1, 1:1, and 1:5, respectively (Liu et al. 2018 ). These acid rain solutions were applied to each plot twice monthly using a sprinkler, from March 2015 to February 2016. The total quantity of applied simulated acid rain was 62.07 mm, the proportions of which were based on the monthly precipitation from 2002 to 2013 (Liu et al. 2017a ). The quantities of simulated acid rain during the fine root litter decomposition periods (3 months, 6 months, 9 months, and 12 months) were 12.79, 45.15, 54.49, and 62.07 mm, respectively.

The S3 (S/N 5:1, pH 2.5), SN3 (1:1, 2.5), and N3 (1:5, 2.5) treatments supplied 0.25 g m −2 , 0.92 g m −2 , 2.29 g m −2 nitrogen (N), and 2.85 g m −2 , 2.09 g m −2 , 1.05 g m −2 sulfur (S), respectively, to the plots in contrast to the CK. The N and S added to the soil via the pH 2.5 treatments were 10 times that of the pH 3.5 AR treatments, and 100 times that of the pH 4.5 AR treatments.

The decomposition of fine root litter was determined using the standard litterbag technique, with the detailed procedure described by Liu et al. ( 2017a , b ). We collected fresh fine roots (first- or second-order roots Ø ≤ 2 mm) in November 2014 (following the growing season) from the plantation soils under study at a depth of 10 cm, and the fine root litter was manually washed free of soil. Prior to the experiment, all of the fine roots from each species were air-dried and the moisture contents were measured (Liu et al. 2017a ). Fine root litterbags (10 cm × 10 cm, 0.1 mm mesh polyethylene) were placed in the soil at a depth of 10 cm (at a 45° angle relative to the soil surface) in each plot using a hoe with 3 g air-dried roots (water content 8.41%) (Liu et al. 2017a ) in late February 2015. The initial water, total carbon (TC), total nitrogen (TN), total phosphorus (TP), potassium (TK), calcium (Ca), and aluminum (Al) contents of the fine root litter were determined (Additional file 1 : Table S1).

Three fine root litterbags from each of the three simulated acid rain treated plots were collected after 3 months (March to May 2015), 6 months (August 2015), 9 months (November 2015), and 12 months (February 2016) (Additional file 1 : Fig. S2). Meanwhile, we collected soil samples at a depth of 10 cm from around the litterbags. The soil samples were sifted through a 2 mm sieve to remove leaves, plant roots, gravel, and stones (Liu et al. 2017a ). Any soil that adhered to the fine root litter samples was carefully removed, and the litter was manually rinsed with distilled water. The fine root litters were oven-dried at 60 °C for 24 h to a constant weight for the determination of mass loss.

Soil and litter properties

The soil pH was determined using a PB-10 pH meter, and the soil enzyme activities were determined via a spectrophotometer as described by Liu et al. ( 2017a ) and (Liu et al. 2020 ), respectively. The fine root C and N were determined using an elemental analyzer (Vario EL III, Elementar, Germany) and described by Liu et al. ( 2018 ). The phosphorus (P) concentration was determined using a UV-Vis spectrophotometer, where a 0.25 g sample was extracted using concentrated nitric acid and perchloric acid (5:1). The digestion solution, which included K, Ca, and Al ions, was quantitatively analyzed using an atomic absorption spectrometer (AA900T, Perkin Elmer, MA, USA).

Statistical analyses

The remaining fine root mass of each sample was expressed as a percentage of the initial fine root dry weight. We applied a first-order exponential decay model X t / X o  = e −kt to fit the decomposition data (Song et al. 2017 ), where X t is the net oven-dried weight remaining at time t ; X o is the initial oven-dried weight; and k is the annual decomposition rate constant (yr −1 ) (Song et al. 2017 ). The quantity of elements released from the decomposing fine roots were expressed as percentages of the initial element content, which were calculated by the Eq. E  = [( M t  ×  C t )/( M o  ×  C o )] × 100, where E is the quantity of remaining elements (%); M t is the oven-dried mass at time t ; and C t is the nutrient concentration at time t (mg g −1 ); M o is the initial oven-dried mass (g); C o is the initial element concentration (mg g −1 ); E  > 100 indicates element immobilization; and E  < 100 indicates element release (Brandt et al. 2010 ; Song et al. 2017 ).

To test whether the soil traits and annual fine root decomposition rates varied under the acid rain treatments, we conducted a series of one-way analyses of variance (ANOVA) with a Duncan test using SPSS 19.0 (SPSS Inc., Chicago, Ill., USA). Three-way ANOVA was employed to assess the effects of the decomposition period, acid rain pH, and S/N ratio on the soil traits and litter decomposition using SPSS 19.0. Pearson analysis was performed to reveal the relationships between the fine root decomposition and soil properties under different acid rain treatments in R 3.4.0. Pearson and structural equation modelling analyses were performed using R 3.4.0. and AMOS 24.0, respectively.

The range of pH values under the CK treatments was 4.04 ± 0.03 ~ 4.32 ± 0.04 (Fig.  1 ). Under acid rain stress, the range of soil pH values shifted to 3.72 ± 0.03 ~ 4.21 ± 0.07. However, no significant differences in the soil pH were observed between the different acid rain S/N ratios ( p  > 0.05) . Furthermore, there was a significant difference between the decomposing periods ( p  < 0.001), with the soil pH values after 6 months being the lowest. The soil pH values were significantly higher for the CK treatment than under simulated acid rain treatments at pH 3.5 and 2.5 ( p  < 0.05), except for those at 6 months.

Change trends of soil pH values under acid rain stress. The experimental treatments were: CK, control check; S1, pH 4.5, S/N 5:1; S2, pH 3.5, S/N 5:1; S3, pH 2.5, S/N 5:1; SN1, pH 4.5, S/N 1:1; SN2, pH 3.5, S/N 1:1; SN3, pH 2.5, S/N 1:1; N1, pH 4.5, S/N 1:5; N2, pH 3.5, S/N 1:5; N3, pH 2.5, S/N 1:5. Different letters indicate significant difference ( p  < 0.05) between different treatments in the same season based on one-way ANOVA, followed by a Duncan test. Three-way ANOVA was applied to indicate significant differences between variances (no CK treatments). S/N, the ratio of SO 4 2− to NO 3 − ; pH, acid rain pH, ***, ** and * indicate significant differences at p  < 0.001, 0.01, and 0.05, respectively

Soil enzyme activities

The variable trends in soil enzyme activities were similar to those of the soil pH over the four decomposition periods (Fig.  2 ), where significant differences were observed between decomposition periods (Additional file 1 : Table S2, p  < 0.001). In addition, the acid rain pH significantly influenced the activities of soil enzymes over the four decomposition periods ( p  < 0.001 or 0.01, Additional file 1 : Table S2). Acid rain decreased the soil enzyme activities in contrasted to the CK treatments (Fig. 2 A, C, D), except for the pH 4.5 acid rain for acid phosphatase activity (Fig.  2 B). Furthermore, the pH 2.5 treatments significantly decreased the soil enzyme activities over the four decomposition periods ( p  < 0.05). However, we found that only the S/N ratio influenced the activities of urease, acid phosphatase, and catalase over 6 months (Additional file 1 : Table S2). The soil urease activity increased significantly with higher acid rain NO 3 − content ( p  < 0.05, Fig.  2 A).

Effects of simulated acid rain on soil enzyme activities over four decomposition periods. Values in brackets are standard deviations ( n  = 3). Other details as in Fig.  1

Fine root mass loss and k value of decomposition rate

The fine root litter masses in the C. lanceolata plantation were rapidly lost during the 1-year experimental period (Fig. 3 A–C). The remaining masses under the sulfuric acid rain (S) treatments (84.04 ± 2.78% ~ 86.22 ± 1.74%) during 3 months of decomposition were lower than those under the CK treatments (86.36 ± 4.62%). However, the remaining masses under the mixed sulfuric and nitric acid rain (SN) treatments were greater at lower pH, whereas the nitric acid rain (N) treatments increased the remaining masses during 3 months of decomposition. The minimum fine root litter mass losses under the S, SN, and N treatments occurred under the pH 2.5 treatments after 1 year of decomposition, which were 8.23%, 10.37%, and 6.09% lower than under the CK treatment, respectively. Furthermore, significant inhibitory effects of acid rain pH on mass loss were observed ( p  < 0.01, Additional file 1 : Table S3).

Percentages of mass remaining and annual decomposition rate ( k value) from decomposing fine roots in a C. lanceolata plantation under acid rain stress. A Sulfuric acid rain; B sulfuric and nitric acid rain mixture; C nitric acid rain; D k value, calculated using the first-order exponential decay model ( X t / X o  = e − kt ). Other details as in Fig.  1

The annual fine root litter decomposition rate ( k value) decreased under lower acid rain pH (Fig.  3 D). The k values under the SN3 and N3 treatments were significantly lower than those under the CK treatment ( p  < 0.05). However, no significant differences were observed between the acid rain types ( p  = 0.871).

Dynamics of TC, TN, and TP during fine root decomposition

The TC, TN, and TP (% of initial) assessed in this study exhibited significant temporal patterns (Fig.  4 , Additional file 1 : Table S3). The fine root TC (% of initial) under all treatments decreased over time (Fig. 4 A–C), where the acid rain pH significantly inhibited the C release of the fine root litter ( p  = 0.004, Additional file 1 : Table S3). The fine root TN (% of initial) under the CK, S, and SN treatments rapidly decreased over 3 months, and then showed no obvious changes at 6 months (except for SN1 and SN2) (Fig. 4 A1, B1). However, the TN under the N1 treatment initially increased and then decreased (Fig. 4 C1). In addition, the acid rain S/N ratio significantly influenced the TN over the four decomposing periods ( p  < 0.001, Additional file 1 : Table S3).

Percentages of TC A , B , and C , TN A1 , B1 , and C1 and TP A2 , B2 , and C1 remaining in the decomposing fine roots of a C. lanceolata plantation in a 1-year fine root litter decomposition experiment under acid rain stress. A , A1 , and A2 : sulfuric acid rain; B , B1 , and B2 : sulfuric and nitric acid rain mixture; C , C1 , and C2 nitric acid rain. Other details as in Fig.  1

The fine root TP (% of initial) under the CK, S, SN1, and SN3 treatments initially increased, and then decreased after 6 months (Fig. 4 A2, B2). However, the SN2 and N treatments were shown to decrease the TP over 3 months (Fig. 4 B2, C2). After 1 year of decomposition, the remaining percentage and content of TP under the CK treatment were 64.12 ± 6.19% and 0.60 ± 0.03 g kg – 1 , respectively. Moreover, the TP under the N treatments (50.48 ± 6.25% ~ 57.19 ± 1.09%) was lower than under the S (62.79 ± 6.25% ~ 64.71 ± 10.49%) and SN treatments (60.34 ± 12.17% ~ 63.76 ± 3.16%). Thus, statistically significant interactions between the influences of acid rain pH and the S/N ratio on the TP were found ( p  < 0.05, Additional file 1 : Table S3).

Metal dynamics during fine root decomposition

The fine root litter TK (% of initial) exhibited a unique pattern in contrast to the TC, TN, and TP in that the TK concentration initially rapidly decreased, which decelerated following 3 months of decomposition (Fig. 5 A–C). ANOVA analysis revealed that the acid rain pH, S/N ratio, and decomposition time had significant effects on the release of K from the fine root litter ( p  < 0.01, Additional file 1 : Table S3).

Percentages of K ( A , B , and C ), Ca ( A1 , B1 , and C1 ), and Al ( A2 , B2 , and C1 ) remaining in the decomposing fine roots of a C. lanceolata plantation in a 1-year fine-root litter decomposition experiment under acid rain stress. A , A1 , and A2 : sulfuric acid rain; B , B 1 and B 2: sulfuric and nitric acid rain mixture; C , C1 , and C2 : nitric acid rain. Other details as in Fig.  1

The fine root Ca was immobilized under the CK treatments over 3 months of decomposition. However, acid rain, except for SN1 and N1, accelerated the release of Ca from the fine root litter (Fig. 5 A1, B1, C1), particularly under the impacts of stronger acid rain (pH = 2.5). Over 6 months of decomposition, the remaining mass of Ca under the CK treatment rapidly decreased. Furthermore, we observed that the remaining mass of Ca following 1 year of decomposition increased under the effects of the S1 treatments.

Under the CK, S1, S3, SN3, and N treatments, the fine root Al was immobilized over 3 months of decomposition, particularly under the N3 treatment (Fig. 5 A2, B2, C2). Furthermore, under the S2 and N3 treatments, the fine root Al (% of initial) increased following 12 months of decomposition (compared to 9 months) and was higher than that under the CK treatment. However, the S1, S3, SN1, and SN2 treatments reduced the fine root Al in contrast to the CK.

ANOVA analysis indicated that the acid rain pH and S/N ratio had significant effects (both separately and in combination) on the fine root Ca and Al of the fine root litter ( p  < 0.001 or 0.01, Additional file 1 : Table S3), except for interactive influences on the fine root Ca ( p  = 0.274, Additional file 1 : Table S3).

Linking fine root decomposition to soil properties and fine root elements

There were significant positive correlations between the decomposition rate and soil pH, soil TN, acid phosphatase activity, and sucrase activity (Additional file 1 : Figs. S3, S4). However, negative correlations were observed between the decomposition rate and fine root C/N ratio, Al, soil C/N ratio, and soil available potassium content (AK), albeit they were not significant (Additional file 1 : Fig. S4). In addition, the acid rain pH had significant positive effects on the activities of soil enzymes ( p  < 0.001) and fine root Ca content ( p  < 0.01, Additional file 1 : Fig. S4). Furthermore, the acid rain S/N ratio had significant positive effects on the fine root Ca content ( p  < 0.05, Additional file 1 : Fig. S4).

The soil and fine root indices were selected according to the correlation analysis between all indices (Additional file 1 : Fig. S4) and the model parameter (Fig.  6 ). The total effect of acid rain pH (0.476) on the fine root decomposition rate was stronger than that of the acid rain S/N ratio (0.090) (Table 1 ). Furthermore, the direct effects of acid rain pH on soil properties (pHs 0.948, phosphatase 0.920), except for soil urease, were stronger than those of the acid rain S/N ratio (pHs 0.134, phosphatase − 0.041) (Fig.  6 ). However, the direct effects of the acid rain S/N ratio on the fine root Ca (0.322) and C/N ratio (− 0.235) were stronger than the impacts on soil pH and soil phosphatase activity. The indirect effects of acid rain pH (0.757) on the fine root decomposition rate proceeded primarily through altered soil pH and enzyme activities. However, the indirect effects of the acid rain S/N ratio (0.265) on the decomposition rate occurred principally through changes in the fine root C/N ratio, Ca content, and soil urease activity (Table 1 ).

Structural equation models of the impacts of acid rain on fine root decomposition rate. (χ 2  = 5.325; p  = 0.324 > 0.05; GFI = 0.952 > 0.900; RMSEA = 0.077 < 0.08). Numbers on arrows are standardized path coefficients. Solid lines indicate the direct influence of each parameter on the decomposition rate and dotted lines indicate the indirect influence of acid rain on the decomposition rate through changing soil or root parameters. pHs, soil pH, phosphatase, soil acid phosphatase activity; urease, soil urease activity; Car, calcium content of fine root litter; C/Nr, C/N ratio of fine root litter; rate, fine root litter decomposition rate; S/N, the ratio of SO 4 2− to NO 3 − ; pH, acid rain pH

Recently, China has brought SO 2 emissions under control toward the mitigation of acid rain contamination. However, NO 2 emissions have not decreased significantly, which has led to progressive changes in the composition of acid rain, from sulfuric acid rain (SAR) to nitric acid rain (NAR) (Liu et al. 2020 ). This study represents an attempt to elucidate how different acid rain treatments impact fine root decomposition patterns through changing abiotic and biotic factors.

Effects of acid rain on fine root decomposition through changing abiotic factors

The decomposition of fine roots is an important resource for the soil C pool, which is primarily influenced by climate, soil environment, and its own quality (Allison et al. 2018 ; Li et al. 2020 ). In this study, we found that the fine root decomposition rate of C. lanceolata decreased under lower acid rain pH, but not significantly (Fig.  3 D), which was consistent with our previous studies regarding the decomposition of Q. acutissima litter (Liu et al. 2017a ). In addition, strongly acidic rain with higher NO 3 − concentrations (SN3 and N3) significantly decreased the fine root decomposition rate compared with the CK treatment (Fig.  3 D), which supported our second hypothesis. Previous investigations revealed that acid rain with an increased NO 3 − content aggravated soil acidification (Liu et al. 2019 , 2017a ) and affected the decomposition of fine roots, which was not significant in the present study (Additional file 1 : Fig. S4). This may have been due to the experimental period (1 year) of this study being too short. In addition, based on SEM analysis, we found that the overall inhibitory effect of acid rain pH on fine root decomposition was stronger than that of acid rain NO 3 − , which was consistent with our previous research results on fine root growth (Liu et al. 2018 ). This may have been due to acid rain pH (0.948) having a stronger effect on soil acidification than the NO 3 − content (0.134) (Fig.  6 ).

In contrast to our findings, several other studies found that acid rain could neutrally or positively impact the decomposition of litter in diverse ecosystems (Lv et al. 2014 ; Wei et al. 2020 ). This may have been due to the differences in plant species, acid rain intensity, and simulation time (Wei et al. 2020 ). Furthermore, the S and N of acid rain can provide nutrients for decomposers, thus stimulating the decomposition of litter, particularly during the initial stages of acid rain simulation (Liu et al. 2017a ).

As nutrients from fine root litter are released into the soil, their basic properties are altered, which are critical driving factors for litter decomposition (Wang et al. 2019 ). Diverse elemental constituents showed variable dynamics under particular compositions of acid rain. A previous study found that acid rain could alter the litter decomposition rate by indirectly affecting litter quality (Wei et al. 2020 ). However, in this study, we found that the acid rain pH had no significant effects on the TC and TN contents of fine root litter over the four decomposition periods (Additional file 1 : Fig. S4), which was consistent with previous studies (Wei et al. 2020 ). Conversely, we found that the S/N ratio of acid rain significantly influenced the remaining TN (% of initial, Additional file 1 : Table S3) and TN contents over 3 and 9 months (Additional file 1 : Table S4). This may have been because the acid rain S/N significantly influences the activities of soil urease (Additional file 1 : Table S2), which impacts the N cycles of ecosystems (Yang et al. 2017 ).

Phosphorus (P) is a critical element that affects the growth and reproduction of microorganisms. The release of P from litter promotes the growth of microorganisms, thereby facilitating the decomposition of litter (Pourhassan et al. 2016 ). In our study, the TP (% of initial) of C. lanceolata fine root litter under the CK treatment increased over 6 months (Fig.  4 ). This might explain why the P limitation in gymnosperm litter was observed only at the onset of the growing season (Jean et al. 2013 ), as the P was not yet being effectively mobilized through the decomposition of litter (Pourhassan et al. 2016 ).

Furthermore, nitric acid rain shortened the duration of P enrichment (3 months, Fig.  4 C2), which indirectly implied that P limitation might appear earlier under nitric acid stress (Jean et al. 2013 ). Prior studies revealed that high-quality litter with higher N and P concentrations and lower C/N ratios decomposed more quickly (Patoine et al. 2017 ). This was consistent with our results, as there was a positive correlation between the TN and TP contents and the decomposition rate, whereas a negative correlation was found between the C/N and decomposition rates (Additional file 1 : Fig. S4).

For this study, we found that the litter content of Ca (1.95 ± 0.27 g kg −1 ) and Al (5.13 ± 0.43 g kg −1 ) (for the CK treatment during the early decomposition stage) were higher than the initial contents Ca (1.66 ± 0.19 g kg −1 ) and Al (4.12 ± 0.29 g kg −1 ) (Additional file 1 : Table S5), which was consistent with multiple earlier studies (Cao et al. 2018 ; Yue et al. 2016 ). This might have been attributed to potential mechanisms such as the formation of chelates and complexes at active sites on the organic molecules of decomposing detritus (Yue et al. 2016 ).

However, there was a significantly positive correlation between the Ca content and the acid rain pH (Additional file 1 : Fig. S4). This indicated that stronger acid rain accelerated the release of Ca from fine root litter. Moreover, McCay et al. ( 2013 ) demonstrated that invertebrate soil dwelling organisms such as earthworms prefer calcium-rich substances for food, which promotes litter decomposition. Thus, we speculated that acid rain reduced the Ca content of fine root litter, which may have decreased the activity of earthworms that indirectly enhanced the inhibition of acid rain on the decomposition of fine roots.

Effects of acid rain on fine root decomposition through changing biotic factors

Soil enzymes play critical roles in the cycling of nutrients in forest ecosystems, such as urease and phosphatase, which participate in the nitrogen and phosphorus cycles (Gu et al. 2019 ; Liu et al. 2020 ). In our study, we observed that the activities of soil enzymes (e.g., urease, acid phosphatase, sucrose, and catalase) significantly decreased under lower acid rain pH (Fig.  2 ), which was consistent with the results of prior studies (Lv et al. 2014 ; Chen et al. 2019 ). This was because acid rain pH had the capacity to directly damage soil enzyme systems (Wei et al. 2020 ), which further influenced the decomposition of litter.

We also found that the soil pH significantly affected the patterns of soil enzyme activities (Additional file 1 : Fig. S4), which indirectly influenced the decomposition of fine root litter, as indicated by earlier studies (Liu et al. 2019 ; Lv et al. 2014 ). Furthermore, the inhibition of litter decomposition can result in the reduction of substrates that are directly utilized by microorganisms, which translates to lower soil enzyme activities (Tang et al. 2019 ). Thus, we found that soil enzyme activities were positively correlated with the decomposition rate of fine roots, which was consistent with our rationale above.

Previous studies revealed that the addition of N had positive, neutral, or negative effects on the activities of soil enzymes (Jing et al. 2017 ; Liu and Zhang, 2019 ). In this study, we found that the activities of soil urease intensified with the higher N content of acid rain (Fig.  4 A), which was consistent with the reports of Guan et al. ( 2019 ). However, the soil urease activities under the acid rain treatments were still lower than that of the CK. This was primarily due to the inhibitory effects of acid rain pH being stronger than promotional effects of the added N. Furthermore, Guan et al. ( 2019 ) found that the addition of N stimulated microbial activities and increased P concentrations, which enhanced the presence of P acquiring enzymes. In this study, the activities of acid phosphatase increased only under the weaker acid rain treatments in contrast to the CK treatment (Fig.  4 B). However, more potent acid rain (pH 2.5) with higher N inputs significantly decreased phosphatase activity. This verified once again that the inhibitory effects of acid rain pH on soil enzyme activities were stronger than the promotional effects of N inputs.

Following a four-season mesocosm experiment, acid rain treatments led to modifications in elemental dynamics and soil properties during the decomposition of fine root litter in the Yangtze River Delta region of China. Both acid rain pH value and NO 3 − content aggravated soil acidification, which inhibited the decomposition of fine roots. The effects of soil enzyme activities on the decomposition of fine root litter were stronger than the elemental dynamics under acid rain stress. The inhibitory effects of acid rain pH, which altered the soil pH and enzyme activities on the decomposition of fine root litter, were greater than the impacts of the acid rain S/N ratio through changes in the fine root C/N ratio, Ca content, and soil urease activities. In summary, we found that acid rain altered the decomposition of fine roots, primarily by modifying the activities of soil enzymes through acid rain pH; however, changes in the S/N ratio exacerbated the negative effects of acid rain pH.

Availability of data and materials

The data sets used and/or analyzed during the current study are available from the corresponding author on reasonable request.

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Acknowledgements

We would like to thank Frank Boehm, from Lakehead University, for the language editing of this manuscript.

This research was funded by the Jiangsu Province Science Foundation for Youths (BK20200785), the China Postdoctoral Science Foundation (2018M642260), the Jiangsu Agriculture Science and Technology Innovation Fund (CX(17)1004), and the Priority Academic Program Development of Jiangsu Higher Education Institutions (PAPD).

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Xin Liu, Miaojing Meng, Chong Li, Shilin Ma, Qinyu Li & Jinchi Zhang

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Jiangxi Academy of Forestry, 1629 Fenglin Road, Nanchang, 330032, Jiangxi, China

Co-Innovation Center for Sustainable Forestry in Southern China, Nanjing Forestry University, 159 Longpan Road, Nanjing, 210037, Jiangsu, China

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JZ and XL conceived and designed the study. MM and YZ revised and perfected the design of the experiments. XL, CL, SM, QL and QR performed the experiments. Xin Liu wrote the paper. YZ and JZ reviewed and edited the manuscript. All authors read and approved the manuscript.

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Supplementary Information

Additional file 1: table s1..

Initial chemistry of C. lanceolata fine roots. Table S2. The p values of variance analysis under the impacts of simulated acid rain on soil enzyme activities. Table S3. The p values of variance analysis under the impacts of acid rain on remaining mass and element contents of fine root litter. Table S4. Contents of biomass, TC, TN, and C/Nr of decomposing C. lanceolata fine-roots under the stress of acid rain. Table S5. Contents of TP, K, Ca and Al of decomposing C. lanceolata fine-roots under the stress of acid rain. Fig. S1. Distribution map of annual mean precipitation pH value of China in 2020 according to the Environment Bulletin of China in 2020. Fig. S2. Schematic diagram of experimental layout. Fig. S3. Effects of simulated acid rain treatments on soil nutrients over four decomposition periods. Fig. S4. The Pearson relationships between fine root decomposition and soil properties under the stress of acid rain.

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Liu, X., Meng, M., Zhang, Y. et al. Effects of sulfuric, nitric, and mixed acid rain on the decomposition of fine root litter in Southern China. Ecol Process 10 , 65 (2021). https://doi.org/10.1186/s13717-021-00334-0

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4.7: Acid Base Reactions

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Learning Objectives

• To know the characteristic properties of acids and bases

Acid–base reactions are essential in both biochemistry and industrial chemistry. Moreover, many of the substances we encounter in our homes, the supermarket, and the pharmacy are acids or bases. For example, aspirin is an acid (acetylsalicylic acid), and antacids are bases. In fact, every amateur chef who has prepared mayonnaise or squeezed a wedge of lemon to marinate a piece of fish has carried out an acid–base reaction. Before we discuss the characteristics of such reactions, let’s first describe some of the properties of acids and bases.

Definitions of Acids and Bases

In Chapter 4.6 , we defined acids as substances that dissolve in water to produce H + ions, whereas bases were defined as substances that dissolve in water to produce OH − ions. In fact, this is only one possible set of definitions. Although the general properties of acids and bases have been known for more than a thousand years, the definitions of acid and base have changed dramatically as scientists have learned more about them. In ancient times, an acid was any substance that had a sour taste (e.g., vinegar or lemon juice), caused consistent color changes in dyes derived from plants (e.g., turning blue litmus paper red), reacted with certain metals to produce hydrogen gas and a solution of a salt containing a metal cation, and dissolved carbonate salts such as limestone (CaCO 3 ) with the evolution of carbon dioxide. In contrast, a base was any substance that had a bitter taste, felt slippery to the touch, and caused color changes in plant dyes that differed diametrically from the changes caused by acids (e.g., turning red litmus paper blue). Although these definitions were useful, they were entirely descriptive.

The Arrhenius Definition of Acids and Bases

The first person to define acids and bases in detail was the Swedish chemist Svante Arrhenius (1859–1927; Nobel Prize in Chemistry, 1903). According to the Arrhenius definition , an acid is a substance like hydrochloric acid that dissolves in water to produce H + ions (protons; Equation \(\PageIndex{1}\) ), and a base is a substance like sodium hydroxide that dissolves in water to produce hydroxide (OH − ) ions (Equation \(\PageIndex{2}\) ):

\[ \underset{an\: Arrhenius\: acid}{HCl_{(g)}} \xrightarrow {H_2 O_{(l)}} H^+_{(aq)} + Cl^-_{(aq)} \]

\[ \underset{an\: Arrhenius\: base}{NaOH_{(s)}} \xrightarrow {H_2O_{(l)}} Na^+_{(aq)} + OH^-_{(aq)} \]

According to Arrhenius, the characteristic properties of acids and bases are due exclusively to the presence of H + and OH − ions, respectively, in solution. Although Arrhenius’s ideas were widely accepted, his definition of acids and bases had two major limitations:

• First, because acids and bases were defined in terms of ions obtained from water, the Arrhenius concept applied only to substances in aqueous solution.
• Second, and more important, the Arrhenius definition predicted that only substances that dissolve in water to produce \(H^+\) and \(OH^−\) ions should exhibit the properties of acids and bases, respectively. For example, according to the Arrhenius definition, the reaction of ammonia (a base) with gaseous HCl (an acid) to give ammonium chloride (Equation \(\PageIndex{3}\) ) is not an acid–base reaction because it does not involve \(H^+\) and \(OH^−\):

\[ NH_{3\;(g)} + HCl_{(g)} \rightarrow NH_4Cl_{(s)} \]

The Brønsted–Lowry Definition of Acids and Bases

Because of the limitations of the Arrhenius definition, a more general definition of acids and bases was needed. One was proposed independently in 1923 by the Danish chemist J. N. Brønsted (1879–1947) and the British chemist T. M. Lowry (1874–1936), who defined acid–base reactions in terms of the transfer of a proton (H + ion) from one substance to another.

According to Brønsted and Lowry, an acid ( A substance with at least one hydrogen atom that can dissociate to form an anion and an \(H^+\) ion (a proton) in aqueous solution, thereby forming an acidic solution) is any substance that can donate a proton, and a base (a substance that produces one or more hydroxide ions (\(OH^-\) and a cation when dissolved in aqueous solution, thereby forming a basic solution) is any substance that can accept a proton. The Brønsted–Lowry definition of an acid is essentially the same as the Arrhenius definition, except that it is not restricted to aqueous solutions. The Brønsted–Lowry definition of a base, however, is far more general because the hydroxide ion is just one of many substances that can accept a proton. Ammonia, for example, reacts with a proton to form \(NH_4^+\), so in Equation \(\PageIndex{3}\), \(NH_3\) is a Brønsted–Lowry base and \(HCl\) is a Brønsted–Lowry acid. Because of its more general nature, the Brønsted–Lowry definition is used throughout this text unless otherwise specified.

Polyprotic Acids

Acids differ in the number of protons they can donate. For example, monoprotic acids (a compound that is capable of donating one proton per molecule) are compounds that are capable of donating a single proton per molecule. Monoprotic acids include HF, HCl, HBr, HI, HNO 3 , and HNO 2 . All carboxylic acids that contain a single −CO 2 H group, such as acetic acid (CH 3 CO 2 H), are monoprotic acids, dissociating to form RCO 2 − and H + ( section 4.6 ). A compound that can donate more than one proton per molecule. can donate more than one proton per molecule. For example, H 2 SO 4 can donate two H + ions in separate steps, so it is a diprotic acid (a compound that can donate two protons per molecule in separate steps) and H 3 PO 4 , which is capable of donating three protons in successive steps, is a triprotic acid (a compound that can donate three protons per molecule in separate steps), (Equation \(\PageIndex{4}\), Equation \(\PageIndex{5}\), and Equation \(\PageIndex{6}\) ):

\[ H_3 PO_4 (l) \overset{H_2 O(l)}{\rightleftharpoons} H ^+ ( a q ) + H_2 PO_4 ^- (aq) \tag{8.7.4}\]

\[ H_2 PO_4 ^- (aq) \rightleftharpoons H ^+ (aq) + HPO_4^{2-} (aq) \tag{8.7.5}\]

\[ HPO_4^{2-} (aq) \rightleftharpoons H^+ (aq) + PO_4^{3-} (aq) \tag{8.7.6}\]

In chemical equations such as these, a double arrow is used to indicate that both the forward and reverse reactions occur simultaneously, so the forward reaction does not go to completion. Instead, the solution contains significant amounts of both reactants and products. Over time, the reaction reaches a state in which the concentration of each species in solution remains constant. The reaction is then said to be in equilibrium (the point at which the rates of the forward and reverse reactions become the same, so that the net composition of the system no longer changes with time).

Strengths of Acids and Bases

We will not discuss the strengths of acids and bases quantitatively until next semester. Qualitatively, however, we can state that strong acids ( An acid that reacts essentially completely with water) to give \(H^+\) and the corresponding anion. react essentially completely with water to give \(H^+\) and the corresponding anion. Similarly, strong bases ( A base that dissociates essentially completely in water) to give \(OH^-\) and the corresponding cation) dissociate essentially completely in water to give \(OH^−\) and the corresponding cation. Strong acids and strong bases are both strong electrolytes. In contrast, only a fraction of the molecules of weak acids ( An acid in which only a fraction of the molecules react with water) to producee \(H^+\) and the corresponding anion. and weak bases (A base in which only a fraction of the molecules react with water to produce \(OH^-\) and the corresponding cation) react with water to produce ions, so weak acids and weak bases are also weak electrolytes. Typically less than 5% of a weak electrolyte dissociates into ions in solution, whereas more than 95% is present in undissociated form.

In practice, only a few strong acids are commonly encountered: HCl, HBr, HI, HNO 3 , HClO 4 , and H 2 SO 4 (H 3 PO 4 is only moderately strong). The most common strong bases are ionic compounds that contain the hydroxide ion as the anion; three examples are NaOH, KOH, and Ca(OH) 2 . Common weak acids include HCN, H 2 S, HF, oxoacids such as HNO 2 and HClO, and carboxylic acids such as acetic acid. The ionization reaction of acetic acid is as follows:

\[ CH_3 CO_2 H(l) \overset{H_2 O(l)}{\rightleftharpoons} H^+ (aq) + CH_3 CO_2^- (aq) \]

Although acetic acid is very soluble in water, almost all of the acetic acid in solution exists in the form of neutral molecules (less than 1% dissociates), as we stated in section 4.1 . Sulfuric acid is unusual in that it is a strong acid when it donates its first proton (Equation \(\PageIndex{8}\) ) but a weak acid when it donates its second proton ( Equation 8.7.9 ) as indicated by the single and double arrows, respectively:

\[ \underset{strong\: acid}{H_2 SO_4 (l)} \xrightarrow {H_2 O(l)} H ^+ (aq) + HSO_4 ^- (aq) \]

\[ \underset{weak\: acid}{HSO_4^- (aq)} \rightleftharpoons H^+ (aq) + SO_4^{2-} (aq) \]

Consequently, an aqueous solution of sulfuric acid contains \(H^+_{(aq)}\) ions and a mixture of \(HSO^-_{4\;(aq)}\) and \(SO^{2−}_{4\;(aq)}\) ions, but no \(H_2SO_4\) molecules.

The most common weak base is ammonia, which reacts with water to form small amounts of hydroxide ion:

\[ NH_3 (g) + H_2 O(l) \rightleftharpoons NH_4^+ (aq) + OH^- (aq) \]

Most of the ammonia (>99%) is present in the form of NH 3 (g). Amines, which are organic analogues of ammonia, are also weak bases, as are ionic compounds that contain anions derived from weak acids (such as S 2− ).

Table \(\PageIndex{1}\) lists some common strong acids and bases. Acids other than the six common strong acids are almost invariably weak acids. The only common strong bases are the hydroxides of the alkali metals and the heavier alkaline earths (Ca, Sr, and Ba); any other bases you encounter are most likely weak. Remember that there is no correlation between solubility and whether a substance is a strong or a weak electrolyte! Many weak acids and bases are extremely soluble in water.

Note the Pattern

There is no correlation between the solubility of a substance and whether it is a strong electrolyte, a weak electrolyte, or a nonelectrolyte.

Table \(\PageIndex{1}\) Common Strong Acids and Bases

Example \(\PageIndex{1}\)

Classify each compound as a strong acid, a weak acid, a strong base, a weak base, or none of these.

• CH 3 CH 2 CO 2 H
• CH 3 CH 2 NH 2

Given: compound

Asked for: acid or base strength

A Determine whether the compound is organic or inorganic.

B If inorganic, determine whether the compound is acidic or basic by the presence of dissociable H + or OH − ions, respectively. If organic, identify the compound as a weak base or a weak acid by the presence of an amine or a carboxylic acid group, respectively. Recall that all polyprotic acids except H 2 SO 4 are weak acids.

• A This compound is propionic acid, which is organic. B It contains a carboxylic acid group analogous to that in acetic acid, so it must be a weak acid.
• A CH 3 OH is methanol, an organic compound that contains the −OH group. B As a covalent compound, it does not dissociate to form the OH − ion. Because it does not contain a carboxylic acid (−CO 2 H) group, methanol also cannot dissociate to form H + (aq) ions. Thus we predict that in aqueous solution methanol is neither an acid nor a base.
• A Sr(OH) 2 is an inorganic compound that contains one Sr 2 + and two OH − ions per formula unit. B We therefore expect it to be a strong base, similar to Ca(OH) 2 .
• A CH 3 CH 2 NH 2 is an amine (ethylamine), an organic compound in which one hydrogen of ammonia has been replaced by an R group. B Consequently, we expect it to behave similarly to ammonia ( Equation 8.6.7 ), reacting with water to produce small amounts of the OH − ion. Ethylamine is therefore a weak base.
• A HBrO 4 is perbromic acid, an inorganic compound. B It is not listed in Table 8.3 as one of the common strong acids, but that does not necessarily mean that it is a weak acid. If you examine the periodic table, you can see that Br lies directly below Cl in group 17. We might therefore expect that HBrO 4 is chemically similar to HClO 4 , a strong acid—and, in fact, it is.

Exercise \(\PageIndex{1}\)

• CH 3 CH 2 CH 2 CO 2 H
• (CH 3 ) 2 NH
• strong base
• strong acid
• none of these; formaldehyde is a neutral molecule

The Hydronium Ion

Because isolated protons are very unstable and hence very reactive, an acid never simply “loses” an H + ion. Instead, the proton is always transferred to another substance, which acts as a base in the Brønsted–Lowry definition. Thus in every acid–base reaction, one species acts as an acid and one species acts as a base. Occasionally, the same substance performs both roles, as you will see later. When a strong acid dissolves in water, the proton that is released is transferred to a water molecule that acts as a proton acceptor or base, as shown for the dissociation of sulfuric acid:

\[ \underset{acid\: (proton\: donor)}{H_2 SO_4 (l)} + \underset{base\: (proton\: acceptor)} {H_2 O(l)} \rightarrow \underset{acid}{H _3 O^+ (aq)} + \underset{base}{HSO_4^- (aq)} \]

Technically, therefore, it is imprecise to describe the dissociation of a strong acid as producing \(H^+_{(aq)}\) ions, as we have been doing. The resulting \(H_3O^+\) ion, called the hydronium ionis a more accurate representation of \(H^+_{(aq)}\). For the sake of brevity, however, in discussing acid dissociation reactions, we often show the product as \(H^+_{(aq)}\) (as in Equation \(\PageIndex{7}\) ) with the understanding that the product is actually the\(H_3O^+ _{(aq)}\) ion.

Conversely, bases that do not contain the hydroxide ion accept a proton from water, so small amounts of OH − are produced, as in the following:

\( \underset{base}{NH_3 (g)} + \underset{acid}{H_2 O(l)} \rightleftharpoons \underset{acid}{NH_4^+ (aq)} + \underset{base}{OH^- (aq)} \)

Again, the double arrow indicates that the reaction does not go to completion but rather reaches a state of equilibrium. In this reaction, water acts as an acid by donating a proton to ammonia, and ammonia acts as a base by accepting a proton from water. Thus water can act as either an acid or a base by donating a proton to a base or by accepting a proton from an acid. Substances that can behave as both an acid and a base are said to be amphoteric When substances can behave as both an acid and a base. .

The products of an acid–base reaction are also an acid and a base. In Equation \(\PageIndex{11}\), for example, the products of the reaction are the hydronium ion, here an acid, and the hydrogen sulfate ion, here a weak base. In Equation \(\PageIndex{12}\), the products are NH 4 + , an acid, and OH − , a base. The product NH 4 + is called the conjugate acid The substance formed when a Brønsted–Lowry base accepts a proton. of the base NH 3 , and the product OH − is called the conjugate base The substance formed when a Brønsted–Lowry acid donates a proton. of the acid H 2 O. Thus all acid–base reactions actually involve two conjugate acid–base pairs An acid and a base that differ by only one hydrogen ion. All acid–base reactions involve two conjugate acid–base pairs, the Brønsted–Lowry acid and the base it forms after donating its proton, and the Brønsted–Lowry base and the acid it forms after accepting a proton. ; in Equation \(\PageIndex{12}\), they are NH 4 + /NH 3 and H 2 O/OH − .

Neutralization Reactions

A neutralization reaction (a chemical reaction in which an acid and a base react in stoichiometric amounts to produce water and a salt) is one in which an acid and a base react in stoichiometric amounts to produce water and a salt (the general term for any ionic substance that does not have OH− as the anion or H+ as the cation) , the general term for any ionic substance that does not have OH − as the anion or H + as the cation. If the base is a metal hydroxide, then the general formula for the reaction of an acid with a base is described as follows: Acid plus base yields water plus salt . For example, the reaction of equimolar amounts of HBr and NaOH to give water and a salt (NaBr) is a neutralization reaction:

\[ \underset{acid}{HBr(aq)} + \underset{base}{NaOH(aq)} \rightarrow \underset{water}{H_2 O(l)} + \underset{salt}{NaBr(aq)} \]

Acid plus base yields water plus salt.

If we write the complete ionic equation for the reaction in Equation \(\PageIndex{13}\), we see that \(Na^+_{(aq)}\) and \(Br^−_{(aq)}\) are spectator ions and are not involved in the reaction:

\[ H^+ (aq) + \cancel{Br^- (aq)} + \cancel{Na^+ (aq)} + OH^- (aq) \rightarrow H_2 O(l) + \cancel{Na^+ (aq)} + \cancel{Br^- (aq)} \]

The overall reaction is therefore simply the combination of H + (aq) and OH − (aq) to produce H 2 O, as shown in the net ionic equation:

\[ H^+(aq) + OH^-(aq) \rightarrow H_2O(l) \)]

The net ionic equation for the reaction of any strong acid with any strong base is identical to Equation \(\PageIndex{15}\).

The strengths of the acid and the base generally determine whether the reaction goes to completion. The reaction of any strong acid with any strong base goes essentially to completion, as does the reaction of a strong acid with a weak base, and a weak acid with a strong base. Examples of the last two are as follows:

\[ \underset{strong\: acid}{HCl(aq)} + \underset{weak\: base}{NH_3 (aq)} \rightarrow \underset{salt}{NH_4 Cl(aq)} \]

\[ \underset{weak\: acid} {CH_3 CO _2 H(aq)} + \underset{strong\: base}{NaOH(aq)} \rightarrow \underset{salt}{CH _3 CO _2 Na(aq)} + H_2 O(l) \]

Sodium acetate is written with the organic component first followed by the cation, as is usual for organic salts. Most reactions of a weak acid with a weak base also go essentially to completion. One example is the reaction of acetic acid with ammonia:

\[ \underset{weak\: acid}{CH _3 CO _2 H(aq)} + \underset{weak\: base}{NH_3 (aq)} \rightarrow \underset{salt}{CH_3 CO_2 NH_4 (aq)} \]

An example of an acid–base reaction that does not go to completion is the reaction of a weak acid or a weak base with water, which is both an extremely weak acid and an extremely weak base. We will discuss these reactions in more detail in Chapter 16

Except for the reaction of a weak acid or a weak base with water, acid–base reactions essentially go to completion.

In some cases, the reaction of an acid with an anion derived from a weak acid (such as HS − ) produces a gas (in this case, H 2 S). Because the gaseous product escapes from solution in the form of bubbles, the reverse reaction cannot occur. Therefore, these reactions tend to be forced, or driven, to completion. Examples include reactions in which an acid is added to ionic compounds that contain the HCO 3 − , CN − , or S 2− anions, all of which are driven to completion (Figure \(\PageIndex{1}\) ):

\[ HCO_3^- (aq) + H^+ (aq) \rightarrow H_2 CO_3 (aq) \]

\[ H_2 CO_3 (aq) \rightarrow CO_2 (g) + H_2 O(l) \]

\[ CN^- (aq) + H^+ (aq) \rightarrow HCN(g) \]

\[ S ^{2-} (aq) + H^+ (aq) \rightarrow HS^- (aq) \]

\[ HS^- (aq) + H^+ (aq) \rightarrow H_2 S(g) \]

Figure \(\PageIndex{1}\) The Reaction of Dilute Aqueous HCl with a Solution of Na 2 CO 3 Note the vigorous formation of gaseous CO 2 .

The reactions in Equation \(\PageIndex{21}\) are responsible for the rotten egg smell that is produced when metal sulfides come in contact with acids.

Example \(\PageIndex{2}\)

Calcium propionate is used to inhibit the growth of molds in foods, tobacco, and some medicines. Write a balanced chemical equation for the reaction of aqueous propionic acid (CH 3 CH 2 CO 2 H) with aqueous calcium hydroxide [Ca(OH) 2 ] to give calcium propionate. Do you expect this reaction to go to completion, making it a feasible method for the preparation of calcium propionate?

Given: reactants and product

Asked for: balanced chemical equation and whether the reaction will go to completion

Write the balanced chemical equation for the reaction of propionic acid with calcium hydroxide. Based on their acid and base strengths, predict whether the reaction will go to completion.

Propionic acid is an organic compound that is a weak acid, and calcium hydroxide is an inorganic compound that is a strong base. The balanced chemical equation is as follows:

\(2CH_3CH_2CO_2H(aq) + Ca(OH)_2(aq) \rightarrow (CH_3CH_2CO_2)_2Ca(aq) + 2H_2O(l)\)

The reaction of a weak acid and a strong base will go to completion, so it is reasonable to prepare calcium propionate by mixing solutions of propionic acid and calcium hydroxide in a 2:1 mole ratio.

Exercise \(\PageIndex{2}\)

Write a balanced chemical equation for the reaction of solid sodium acetate with dilute sulfuric acid to give sodium sulfate.

\(2CH_3CO_2Na(s) + H_2SO_4(aq) \rightarrow Na_2SO_4(aq) + 2CH_3CO_2H(aq)\)

Stomach acid. An antacid tablet reacts with 0.1 M HCl (the approximate concentration found in the human stomach).

One of the most familiar and most heavily advertised applications of acid–base chemistry is antacids , which are bases that neutralize stomach acid. The human stomach contains an approximately 0.1 M solution of hydrochloric acid that helps digest foods. If the protective lining of the stomach breaks down, this acid can attack the stomach tissue, resulting in the formation of an ulcer . Because one factor that is believed to contribute to the formation of stomach ulcers is the production of excess acid in the stomach, many individuals routinely consume large quantities of antacids. The active ingredients in antacids include sodium bicarbonate and potassium bicarbonate (NaHCO 3 and KHCO 3 ; Alka-Seltzer); a mixture of magnesium hydroxide and aluminum hydroxide [Mg(OH) 2 and Al(OH) 3 ; Maalox, Mylanta]; calcium carbonate (CaCO 3 ; Tums); and a complex salt, dihydroxyaluminum sodium carbonate [NaAl(OH) 2 CO 3 ; original Rolaids]. Each has certain advantages and disadvantages. For example, Mg(OH) 2 is a powerful laxative (it is the active ingredient in milk of magnesia), whereas Al(OH) 3 causes constipation. When mixed, each tends to counteract the unwanted effects of the other. Although all antacids contain both an anionic base (OH − , CO 3 2− , or HCO 3 − ) and an appropriate cation, they differ substantially in the amount of active ingredient in a given mass of product.

Example \(\PageIndex{3}\)

Assume that the stomach of someone suffering from acid indigestion contains 75 mL of 0.20 M HCl. How many Tums tablets are required to neutralize 90% of the stomach acid, if each tablet contains 500 mg of CaCO 3 ? (Neutralizing all of the stomach acid is not desirable because that would completely shut down digestion.)

Given: volume and molarity of acid and mass of base in an antacid tablet

Asked for: number of tablets required for 90% neutralization

A Write the balanced chemical equation for the reaction and then decide whether the reaction will go to completion.

B Calculate the number of moles of acid present. Multiply the number of moles by the percentage to obtain the quantity of acid that must be neutralized. Using mole ratios, calculate the number of moles of base required to neutralize the acid.

C Calculate the number of moles of base contained in one tablet by dividing the mass of base by the corresponding molar mass. Calculate the number of tablets required by dividing the moles of base by the moles contained in one tablet.

A We first write the balanced chemical equation for the reaction:

\(2HCl(aq) + CaCO_3(s) \rightarrow CaCl_2(aq) + H_2CO_3(aq)\)

Each carbonate ion can react with 2 mol of H + to produce H 2 CO 3 , which rapidly decomposes to H 2 O and CO 2 . Because HCl is a strong acid and CO 3 2− is a weak base, the reaction will go to completion.

B Next we need to determine the number of moles of HCl present:

\( 75\: \cancel{mL} \left( \dfrac{1\: \cancel{L}} {1000\: \cancel{mL}} \right) \left( \dfrac{0 .20\: mol\: HCl} {\cancel{L}} \right) = 0. 015\: mol\: HCl \)

Because we want to neutralize only 90% of the acid present, we multiply the number of moles of HCl by 0.90:

\((0.015\: mol\: HCl)(0.90) = 0.014\: mol\: HCl\)

We know from the stoichiometry of the reaction that each mole of CaCO 3 reacts with 2 mol of HCl, so we need

\( moles\: CaCO_3 = 0 .014\: \cancel{mol\: HCl} \left( \dfrac{1\: mol\: CaCO_3}{2\: \cancel{mol\: HCl}} \right) = 0 .0070\: mol\: CaCO_3 \)

C Each Tums tablet contains

\( \left( \dfrac{500\: \cancel{mg\: CaCO_3}} {1\: Tums\: tablet} \right) \left( \dfrac{1\: \cancel{g}} {1000\: \cancel{mg\: CaCO_3}} \right) \left( \dfrac{1\: mol\: CaCO_3} {100 .1\: \cancel{g}} \right) = 0 .00500\: mol\: CaCO_ 3 \)

Thus we need \(\dfrac{0.0070\: \cancel{mol\: CaCO_3}}{0.00500\: \cancel{mol\: CaCO_3}}= 1.4\) Tums tablets.

Exercise \(\PageIndex{3}\)

Assume that as a result of overeating, a person’s stomach contains 300 mL of 0.25 M HCl. How many Rolaids tablets must be consumed to neutralize 95% of the acid, if each tablet contains 400 mg of NaAl(OH) 2 CO 3 ? The neutralization reaction can be written as follows:

\( NaAl(OH)_2CO_3(s) + 4HCl(aq) \rightarrow AlCl_3(aq) + NaCl(aq) + CO_2(g) + 3H_2O(l) \)

6.4 tablets

The pH Scale

One of the key factors affecting reactions that occur in dilute solutions of acids and bases is the concentration of H + and OH − ions. The pH scale A logarithmic scale used to express the hydrogen ion (H + ) concentration of a solution, making it possible to describe acidity or basicity quantitatively. provides a convenient way of expressing the hydrogen ion (H + ) concentration of a solution and enables us to describe acidity or basicity in quantitative terms.

Pure liquid water contains extremely low but measurable concentrations of H 3 O + (aq) and OH − (aq) ions produced via an autoionization reaction , in which water acts simultaneously as an acid and as a base:

\[H_2O(l) + H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)\tag{8.7.22}\)

The concentration of hydrogen ions in pure water is only 1.0 × 10 −7 M at 25°C. Because the autoionization reaction produces both a proton and a hydroxide ion, the OH − concentration in pure water is also 1.0 × 10 −7 M. Pure water is a neutral solution A solution in which the total positive charge from all the cations is matched by an identical total negative charge from all the anions. , in which [H + ] = [OH − ] = 1.0 × 10 −7 M.

The pH scale describes the hydrogen ion concentration of a solution in a way that avoids the use of exponential notation; pH The negative base-10 logarithm of the hydrogen ion concentration: pH=-log[H + ] is defined as the negative base-10 logarithm of the hydrogen ion concentration: pH is actually defined as the negative base-10 logarithm of hydrogen ion activity . As you will learn in a more advanced course, the activity of a substance in solution is related to its concentration. For dilute solutions such as those we are discussing, the activity and the concentration are approximately the same.

\[ pH = -log[H^+]\]

Conversely,

\([ [H^+] = 10^{-pH}\]

Because the hydrogen ion concentration is 1.0 × 10 −7 M in pure water at 25°C, the pH of pure liquid water (and, by extension, of any neutral solution) is

\[ pH = -log[1.0 \times 10^{-7}] = 7.00\]

Adding an acid to pure water increases the hydrogen ion concentration and decreases the hydroxide ion concentration because a neutralization reaction occurs, such as that shown in Equation 8.7.15 . Because the negative exponent of [H + ] becomes smaller as [H + ] increases, the pH decreases with increasing [H + ]. For example, a 1.0 M solution of a strong monoprotic acid such as HCl or HNO 3 has a pH of 0.00:

\[ pH = -log[1.0] = 0.00\]

pH decreases with increasing [H + ].

Conversely, adding a base to pure water increases the hydroxide ion concentration and decreases the hydrogen ion concentration. Because the autoionization reaction of water does not go to completion, neither does the neutralization reaction. Even a strongly basic solution contains a detectable amount of H + ions. For example, a 1.0 M OH − solution has [H + ] = 1.0 × 10 −14 M. The pH of a 1.0 M NaOH solution is therefore

\[ pH = -log[1.0 \times 10^{-14}] = 14.00\]

For practical purposes, the pH scale runs from pH = 0 (corresponding to 1 M H + ) to pH 14 (corresponding to 1 M OH − ), although pH values less than 0 or greater than 14 are possible.

We can summarize the relationships between acidity, basicity, and pH as follows:

• If pH = 7.0, the solution is neutral.
• If pH < 7.0, the solution is acidic.
• If pH > 7.0, the solution is basic.

Keep in mind that the pH scale is logarithmic, so a change of 1.0 in the pH of a solution corresponds to a tenfold change in the hydrogen ion concentration. The foods and consumer products we encounter daily represent a wide range of pH values, as shown in Figure 8.7.2 .

Figure 8.7.2 A Plot of pH versus [H + ] for Some Common Aqueous Solutions

Although many substances exist in a range of pH values (indicated in parentheses), they are plotted using typical values.

Example \(\PageIndex{4}\)

• What is the pH of a 2.1 × 10 −2 M aqueous solution of HClO 4 ?
• The pH of a vinegar sample is 3.80. What is its hydrogen ion concentration?

Given: molarity of acid or pH

Asked for: pH or [H + ]

Using the balanced chemical equation for the acid dissociation reaction and Equation \(\PageIndex{24}\) or \(\PageIndex{25}\), determine [H + ] and convert it to pH or vice versa.

\(HClO_4(l) \rightarrow H^+(aq) + ClO_4^-(aq)\)

The H + ion concentration is therefore the same as the perchloric acid concentration. The pH of the perchloric acid solution is thus

\(pH = -log[H^+] = -log(2.1 \times 10^{-2}) = 1.68\)

The result makes sense: the H + ion concentration is between 10 −1 M and 10 −2 M, so the pH must be between 1 and 2.

Note : The assumption that [H + ] is the same as the concentration of the acid is valid for only strong acids. Because weak acids do not dissociate completely in aqueous solution, a more complex procedure is needed to calculate the pH of their solutions.

\(10^{-pH} = [H^+]\)

Thus \([H^+] = 10^{-3.80} = 1.6 \times 10^{-4}\: M\).

Exercise \(\PageIndex{4}\)

• What is the pH of a 3.0 × 10 −5 M aqueous solution of HNO 3 ?
• What is the hydrogen ion concentration of turnip juice, which has a pH of 5.41?
• \(pH = 4.52\)
• \([H^+] = 3.9 \times 10^{-6}\: M\)

Tools have been developed that make the measurement of pH simple and convenient ( Figure 8.6.3 ). For example, pH paper consists of strips of paper impregnated with one or more acid–base indicators An intensely colored organic molecule whose color changes dramatically depending on the pH of the solution. , which are intensely colored organic molecules whose colors change dramatically depending on the pH of the solution. Placing a drop of a solution on a strip of pH paper and comparing its color with standards give the solution’s approximate pH. A more accurate tool, the pH meter, uses a glass electrode, a device whose voltage depends on the H + ion concentration.

Figure 8.6.3 Two Ways of Measuring the pH of a Solution: pH Paper and a pH Meter

Note that both show that the pH is 1.7, but the pH meter gives a more precise value.

Key Equations

definition of pH

Equation \(\PageIndex{231}\) : \(pH = -log[H^+]\)

Equation \(\PageIndex{24}\) : \([H^+] = 10^{-pH}\)

Acid–base reactions require both an acid and a base. In Brønsted–Lowry terms, an acid is a substance that can donate a proton (H + ), and a base is a substance that can accept a proton. All acid–base reactions contain two acid–base pairs: the reactants and the products. Acids can donate one proton ( monoprotic acids ), two protons ( diprotic acids ), or three protons ( triprotic acids ). Compounds that are capable of donating more than one proton are generally called polyprotic acids . Acids also differ in their tendency to donate a proton, a measure of their acid strength. Strong acids react completely with water to produce H 3 O + (aq) (the hydronium ion ), whereas weak acids dissociate only partially in water. Conversely, strong bases react completely with water to produce the hydroxide ion, whereas weak bases react only partially with water to form hydroxide ions. The reaction of a strong acid with a strong base is a neutralization reaction , which produces water plus a salt .

The acidity or basicity of an aqueous solution is described quantitatively using the pH scale . The pH of a solution is the negative logarithm of the H + ion concentration and typically ranges from 0 for strongly acidic solutions to 14 for strongly basic ones. Because of the autoionization reaction of water, which produces small amounts of hydronium ions and hydroxide ions, a neutral solution of water contains 1 × 10 −7 M H + ions and has a pH of 7.0. An indicator is an intensely colored organic substance whose color is pH dependent; it is used to determine the pH of a solution.

Key Takeaway

• An acidic solution and a basic solution react together in a neutralization reaction that also forms a salt.

Conceptual Problems

Why was it necessary to expand on the Arrhenius definition of an acid and a base? What specific point does the Brønsted–Lowry definition address?

State whether each compound is an acid, a base, or a salt.

Classify each compound as a strong acid, a weak acid, a strong base, or a weak base in aqueous solution.

• sodium hydroxide
• acetic acid
• magnesium hydroxide
• tartaric acid
• sulfuric acid
• hydroxylamine (NH 2 OH)
• hydrocyanic acid

Decide whether each compound forms an aqueous solution that is strongly acidic, weakly acidic, strongly basic, or weakly basic.

• propanoic acid
• hydrobromic acid
• methylamine
• lithium hydroxide
• citric acid
• sodium acetate
• ammonium chloride
• barium hydroxide

What is the relationship between the strength of an acid and the strength of the conjugate base derived from that acid? Would you expect the CH 3 CO 2 − ion to be a strong base or a weak base? Why? Is the hydronium ion a strong acid or a weak acid? Explain your answer.

What are the products of an acid–base reaction? Under what circumstances is one of the products a gas?

Explain how an aqueous solution that is strongly basic can have a pH, which is a measure of the acidity of a solution.

• weakly acidic
• strongly acidic
• weakly basic
• strongly basic

Numerical Problems

Please be sure you are familiar with the topics discussed in Essential Skills 3 ( section 4.11" )before proceeding to the Numerical Problems.

Derive an equation to relate the hydrogen ion concentration to the molarity of a solution of a strong monoprotic acid.

Derive an equation to relate the hydroxide ion concentration to the molarity of a solution of

• a group I hydroxide.
• a group II hydroxide.

Given the following salts, identify the acid and the base in the neutralization reactions and then write the complete ionic equation:

• barium sulfate
• lithium nitrate
• sodium bromide
• calcium perchlorate

What is the pH of each solution?

• 5.8 × 10 −3 mol of HNO 3 in 257 mL of water
• 0.0079 mol of HI in 750 mL of water
• 0.011 mol of HClO 4 in 500 mL of water
• 0.257 mol of HBr in 5.00 L of water

What is the hydrogen ion concentration of each substance in the indicated pH range?

• black coffee (pH 5.10)
• milk (pH 6.30–7.60)
• tomatoes (pH 4.00–4.40)
• orange juice (pH 3–4)
• fresh egg white (pH 7.60–7.80)
• lemon juice (pH 2.20–2.40)

What is the pH of a solution prepared by diluting 25.00 mL of 0.879 M HCl to a volume of 555 mL?

Vinegar is primarily an aqueous solution of acetic acid. Commercial vinegar typically contains 5.0 g of acetic acid in 95.0 g of water. What is the concentration of commercial vinegar? If only 3.1% of the acetic acid dissociates to CH 3 CO 2 − and H + , what is the pH of the solution? (Assume the density of the solution is 1.00 g/mL.)

If a typical household cleanser is 0.50 M in strong base, what volume of 0.998 M strong monoprotic acid is needed to neutralize 50.0 mL of the cleanser?

A 25.00 mL sample of a 0.9005 M solution of HCl is diluted to 500.0 mL. What is the molarity of the final solution? How many milliliters of 0.223 M NaOH are needed to neutralize 25.00 mL of this final solution?

If 20.0 mL of 0.10 M NaOH are needed to neutralize 15.0 mL of gastric fluid, what is the molarity of HCl in the fluid? (Assume all the acidity is due to the presence of HCl.) What other base might be used instead of NaOH?

Malonic acid (C 3 H 4 O 4 ) is a diprotic acid used in the manufacture of barbiturates. How many grams of malonic acid are in a 25.00 mL sample that requires 32.68 mL of 1.124 M KOH for complete neutralization to occur? Malonic acid is a dicarboxylic acid; propose a structure for malonic acid.

Describe how you would prepare 500 mL of a 1.00 M stock solution of HCl from an HCl solution that is 12.11 M. Using your stock solution, how would you prepare 500 mL of a solution that is 0.012 M in HCl?

Given a stock solution that is 8.52 M in HBr, describe how you would prepare a 500 mL solution with each concentration.

• 4.00 × 10 −3 M

How many moles of solute are contained in each?

• 25.00 mL of 1.86 M NaOH
• 50.00 mL of 0.0898 M HCl
• 13.89 mL of 0.102 M HBr

A chemist needed a solution that was approximately 0.5 M in HCl but could measure only 10.00 mL samples into a 50.00 mL volumetric flask. Propose a method for preparing the solution. (Assume that concentrated HCl is 12.0 M.)

Write the balanced chemical equation for each reaction.

• perchloric acid with potassium hydroxide
• nitric acid with calcium hydroxide
• solid strontium hydroxide with hydrobromic acid
• aqueous sulfuric acid with solid sodium hydroxide

A neutralization reaction gives calcium nitrate as one of the two products. Identify the acid and the base in this reaction. What is the second product? If the product had been cesium iodide, what would have been the acid and the base? What is the complete ionic equation for each reaction?

[H 3 O + ] = [HA] M

• H 2 SO 4 and Ba(OH) 2 ; 2H + + SO 4 2− + Ba 2 + + 2OH − → 2H 2 O + Ba 2 + + SO 4 2−
• HNO 3 and LiOH; H + + NO 3 − + Li + + OH − → H 2 O + Li + + NO 3 −
• HBr and NaOH; H + + Br − + Na + + OH − → H 2 O + Na + + Br −
• HClO 4 and Ca(OH) 2 ; 2H + + 2ClO 4 − + Ca 2 + + 2OH − → 2H 2 O + Ca 2 + + 2ClO 4 −
• 7.9 × 10 −6 M H +
• 5.0 × 10 −7 to 2.5 × 10 −8 M H +
• 1.0 × 10 −4 to 4.0 × 10 −5 M H +

0.13 M HCl; magnesium carbonate, MgCO 3 , or aluminum hydroxide, Al(OH) 3

1.00 M solution: dilute 41.20 mL of the concentrated solution to a final volume of 500 mL. 0.012 M solution: dilute 12.0 mL of the 1.00 M stock solution to a final volume of 500 mL.

• 4.65 × 10 −2 mol NaOH
• 4.49 × 10 −3 mol HCl
• 1.42 × 10 −3 mol HBr
• HClO 4 + KOH → KClO 4 + H 2 O
• 2HNO 3 + Ca(OH) 2 → Ca(NO 3 ) 2 + 2H 2 O

The acid is nitric acid, and the base is calcium hydroxide. The other product is water.

\(2HNO_3 + Ca(OH)_2 \rightarrow Ca(NO_3)_2 + 2H_2O\)

The acid is hydroiodic acid, and the base is cesium hydroxide. The other product is water.

\( HI + CsOH \rightarrow CsI + H_2O \)

The complete ionic equations are

\( 2H^+ + 2NO_3^- + Ca^{2+} + 2OH^- \rightarrow Ca^{2+} + 2NO_3^- + H_2O\) \( H^+ + I^- + Cs^+ + OH^- \rightarrow Cs^+ + I^- + H_2O \)

Contributors

Modified by Joshua Halpern ( Howard University )

Video from Colin McKay @ YouTube