1.7.4

Balancing Equations & Ionic Equations

Test yourself on Balancing Equations & Ionic Equations

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Balanced Equations

Balanced equations are absolutely crucial to any kind of quantitative chemistry.

What are they?

What are they?

  • There are a few key features of balanced equations:
    • They must have an equal number of atoms on either side.
    • You add more atoms by adding reactants or products to either side.
      • There's no method to do this apart from trial and error - just find what works.
  • Remember atoms are rearranged in chemical reactions, so there needs to be an equal number of atoms of each element either side of the equation.
Why are they useful?

Why are they useful?

  • Balanced equations can be used to calculate masses of reactants and volumes of gases.
    • We'll demonstrate these via some worked examples in this session.
Titration

Titration

  • Titration is an experimental technique for finding the concentrations of solutions.
    • Titrations use balanced equations to calculate the concentration of an unknown solution from a volume of a known solution.
    • The experimental technique involves the use of a burette, and often an indicator.
    • A small volume of solution is dropped from the burette into the reaction flask. This is repeated until the indicator shows no further reaction happening.
Molecular equations

Molecular equations

  • When aqueous solutions of CaCl2 and AgNO3 are mixed, a reaction takes place producing aqueous Ca(NO3)2 and solid AgCl:
  • CaCl2(aq)+2AgNO3(aq)⟶Ca(NO3)2(aq) +2AgCl(s)
  • This balanced equation, derived in the usual fashion, is called a molecular equation
    • This is because it doesn’t explicitly represent the ionic species that are present in solution.

State Symbols

State symbols are letters that tell us what state of matter a substance is in. There are four different state symbols:

(s)

(s)

  • (s) is used for solids.
    • E.g. Mg(s)
    • E.g. Na(s)
(l)

(l)

  • (l) is used for pure liquids.
    • E.g. H2O(l)
    • E.g. Hg(l) (mercury is liquid at room temperature).
(aq)

(aq)

  • (aq) is used for aqueous solutions (substances dissolved in water), and therefore it is used more often than (l).
    • E.g. NaOH(aq)
    • E.g. H2SO4(aq)
(g)

(g)

  • (g) is used for gases.
    • E.g. H2(g)
    • E.g. CO2(g)
Why are they useful?

Why are they useful?

  • State symbols are useful to chemists because the state a reactant is produced in can affect experimental techniques.
    • For example, you wouldn't want to carry out a reaction that generates 30 moles of CO2(g) in a 10cm3 sealed vessel - it would explode.
    • If you left your state symbols out, you might try this.

Ionic Equations

Ionic equations are used to describe reactions in solution.

Ionic solids

Ionic solids

  • When ionic compounds dissolve in water, they may dissociate into their constituent ions.
  • These ions are subsequently dispersed homogenously throughout the resulting solution.
  • Ionic compounds dissolved in water are, therefore, more realistically represented as dissociated ions, in this case:
    • E.g. NaCl(aq) ⇌ Na+(aq) + Cl−(aq)
Ionic equations

Ionic equations

  • Net ionic equations are equations that have excluded all the non-reacting spectator ions.
    • The charge on either side of the ionic equation should be equal.
  • Complete ionic equations show spectator ions.
Complete ionic equations

Complete ionic equations

  • Explicitly representing all dissolved ions results in a complete ionic equation.

    • The formulas for the dissolved ionic compounds are replaced by formulas for their dissociated ions:

  • Ca2+(aq) +2Cl-(aq)+ 2Ag+(aq)+2NO3−(aq)⟶Ca2+(aq) +2NO3−(aq)+2AgCl(s)

Spectator ions

Spectator ions

  • Examining the previous equation shows that two chemical species are present in identical form on both sides of the arrow:
  • Ca2+(aq) and NO3−(aq).
    • These are spectator ions.
  • Spectator ions presence is required to maintain charge neutrality.
  • Spectator ions are neither chemically nor physically changed by a reaction.
    • So, they may be eliminated from the equation to yield a more succinct representation called a net ionic equation.
Net ionic equations

Net ionic equations

  • The net ionic equation for the above reaction is:

    • Cl−(aq)+Ag+(aq)⟶AgCl(s)

  • This net ionic equation indicates that solid silver chloride may be produced from dissolved chloride and silver(I) ions.

  • This is regardless of the source of these ions.
    • Molecular and complete ionic equations provide additional information, namely, the ionic compounds used as sources of Cl- and Ag+.
Jump to other topics
1

Structure - Models of the Particulate of Matter

1.1

Introduction to the Particulate Model of Matter

1.1.1States of Matter1.1.2Kinetic Molecular Theory1.1.3Simple Sphere1.1.4Changing State1.1.5Forces Between Particles1.1.6Changes of State1.1.7Predicting States1.1.8Changing State HyperLearning1.1.9Ionization Energy
1.2

The Nuclear Atom

1.2.1Fundamental Particles1.2.2Isotopes & Mass Number1.2.3Mass Spectrometry1.2.4Mass Spectrometry1.2.5Mass Spectroscopy - Analysis1.2.6Extended Response - Mass Spectrometry1.2.7Radioactivity
1.3

Electron Configuration

1.3.1Electron Shells, Sub-Shells & Orbitals1.3.2Electron Configuration1.3.3Absorption & Emission Spectra1.3.4IB Multiple Choice - Atomic Structure1.3.5Extended Response - The Atom & Electrons
1.4

Counting Particles by Mass: The Mole

1.4.1Relative Masses1.4.2The Mole1.4.3IB Multiple Choice - Moles
1.5

Ideal Gases

1.5.1Ideal Gas Equation1.5.2Gas Laws1.5.3IB Multiple Choice - Gases & Kinetics
1.6

Elements, Compounds & Mixtures

1.6.1Fundamental Particles1.6.2Pure Substances1.6.3Purity & The Law of Definite Proportion1.6.4Mixtures1.6.5Separating Mixtures1.6.6IB Multiple Choice - Pure Substances1.6.7Extended Response - Moles & Pure Substances
1.7

States of Matter & Changes of State

1.7.1States of Matter1.7.2Particle Diagrams1.7.3IB Multiple Choice - Solids, Liquids, & Gases1.7.4Balancing Equations & Ionic Equations1.7.5Empirical & Molecular Formulas
1.8

Reacting Masses &. Volumes

1.8.1Balancing Equations & Ionic Equations1.8.2Percentage Yield
1.9

Solutions

1.9.1Molarity1.9.2Particulate Model of Solutions
2

Structure - Models of Bonding & Structure

2.1

The Ionic Model

2.1.1Ionic Bonding
2.2

The Covalent Model

2.2.1Covalent Bonds2.2.2Complex Ions2.2.3Factors Affecting Covalent Bond Strength2.2.4Polarity
2.3

Covalent Structures

2.3.1Ionic Lewis Structures2.3.2Covalent Lewis Structures2.3.3VSEPR Theory2.3.4Resonance Structures2.3.5Allotropes of Carbon2.3.6Formal Charge2.3.7End of Topic Quiz - Bonding2.3.8IB Multiple Choice - Bonding & Lewis Structures2.3.9IB Multiple Choice - Resonance & Formal Charge
2.4

The Metallic Model

2.4.1Metallic Bonding
2.5

From Models to Materials

2.5.1Alloys2.5.2Properties of Polymers2.5.3Polymerisation Reactions2.5.4Introduction to Alkenes2.5.5Reactions of Alkenes2.5.6End of Topic Quiz - Alkenes
2.6

Valence Electrons & Ionic Compounds

2.6.1Valence Electrons & Forming Ions2.6.2Ionic Compounds2.6.3End of Topic Quiz - Atoms & Ions2.6.4IB Multiple Choice - Valence & Ions
2.7

Molecular Shape

2.7.1VSEPR2.7.2Molecular Shapes2.7.3Bond Hybridization2.7.4IB Multiple Choice - VSEPR & Bond Hybridization2.7.5Extended Response - Bonding & Molecular Shape
2.8

Intermolecular Forces

2.8.1Introduction to Intermolecular Forces2.8.2London Dispersion Forces2.8.3Dipole-Dipole Interactions2.8.4Ion-Dipole Interactions2.8.5Hydrogen Bonding2.8.6Large Biomolecules2.8.7IB Multiple Choice - Intermolecular Forces2.8.8Extended Response - Hydrogen Bonding
3

Structure - Classification of Matter

3.1

The Periodic Table: Classification of Elements

3.1.1The Periodic Table3.1.2Trends in the Periodic Table3.1.3End of Topic Quiz - Periodicity & Trends3.1.4Periodic Trends - Electronegativity & Affinity3.1.5End of Topic Quiz - Periodic Trends3.1.6IB Multiple Choice - Periodic Trends3.1.7Extended Response - Periodic Table
3.2

Periodic Trends

3.2.1Atomic & Ionic Radii3.2.2Ionisation & Electronegativity3.2.3Metallic Behaviour & Oxides3.2.4Period 3 Elements & Oxides3.2.5Period 3 Oxides3.2.6End of Topic Quiz - Period 33.2.7Periodic Trends: Blocks3.2.8Redox Chemistry
3.3

Group 1 Alkali Metals

3.3.1Alkali Metals
3.4

Halogens

3.4.1Introduction to Halogens3.4.2Reactions of Halogens3.4.3Chlorine & Chlorate(I)3.4.4End of Topic Quiz - Group 73.4.5Exam-Style Question - Group 73.4.6End of Topic Quiz - Periodicity, Group 2 & 7
3.5

Noble gases, group 18

3.5.1Noble Gases
3.6

Functional Groups: Classification of Organic

3.6.1Nomenclature3.6.2Trends in Physical Properties3.6.3Saturated & Unsaturated Hydrocarbons3.6.4Nomenclature contd.3.6.5Types of Formulas3.6.6Isomerism3.6.7Formulae3.6.8Homologous Series3.6.9Structural Isomers
3.7

Functional Group Chemistry

3.7.1Degrees of Substitution3.7.2Arenes
3.8

Alkanes

3.8.1Alkanes3.8.2Fractional Distillation3.8.3Cracking3.8.4Combustion3.8.5Chlorination3.8.6End of Topic Quiz - Alkanes
3.9

Alcohols

3.9.1Production of Alcohols3.9.2Oxidation of Alcohols3.9.3Oxidation of Alcohols - 23.9.4Elimination Reactions3.9.5End of Topic Quiz - Alcohols3.9.6Exam-Style Question - Alcohols
3.10

Halogenoalkanes

3.10.1Substitution Reactions3.10.2Elimination Reactions3.10.3Ozone Depletion3.10.4End of Topic Quiz - Halogenoalkanes
4

Reactivity - What Drives Chemical Reaction?

4.1

Endothermic & Exothermic Reactions

4.1.1Endothermic & Exothermic Reactions4.1.2Energy Diagrams4.1.3Heat Transfer & Thermal Equilibrium4.1.4IB Multiple Choice - Enthalpy & Energy
4.2

Enthalpy of Reaction, Formation, & Hess' Law

4.2.1Enthalpy of Reaction4.2.2Bond Enthalpies4.2.3Enthalpy of Formation4.2.4Hess’ Law4.2.5Thermodynamics & Kinetic Control4.2.6End of Topic Quiz - Thermodynamics4.2.7IB Multiple Choice - Enthalpy4.2.8Extended Response - Enthalpy Change
4.3

Entropy

4.3.1Introduction to Entropy4.3.2Absolute Entropy & Entropy Change4.3.3Free Energy4.3.4Thermodynamically Favorable Reactions4.3.5IB Multiple Choice - Entropy & Free Energy4.3.6Extended Response - Free Energy
5

Reactivity - How Much, How Fast & How Far?

5.1

Kinetics

5.1.1Collision Theory & Rates of Reaction5.1.2Maxwell-Boltzmann Distribution5.1.3Increasing Rates5.1.4End of Topic Quiz - Kinetics
5.2

Rates of Reaction

5.2.1Rate Equations5.2.2Rate Equations 25.2.3The Effect of a Catalyst
5.3

Stoichometry

5.3.1Stoichiometry
5.4

Le Châtelier’s Principle

5.4.1Introduction to Le Châtelier’s Principle5.4.2Reaction Quotient & Le Châtelier’s Principle5.4.3Temperature & Le Châtelier’s Principle5.4.4IB Multiple Choice -Le Châtelier’s Principle5.4.5Extended Response - Le Châtelier
5.5

Introduction to Equilibrium

5.5.1Equilibrium & Equilibrium Graphs5.5.2Dynamic Equilibria5.5.3Reversible Reactions & Directions
5.6

Equilibrium Constant

5.6.1Calculating the Equilibrium Constant5.6.2Magnitude of the Equilibrium Constant5.6.3Properties of the Equilibrium Constant5.6.4Calculating Equilibrium Concentrations5.6.5Representations of Equilibrium5.6.6IB Multiple Choice - Equilibrium Constants5.6.7Extended Response - Calculating Kc
5.7

Reaction Quotient & Equilibrium Constant

5.7.1Reaction Quotient5.7.2Equilibrium Constants5.7.3Partial Pressures5.7.4Kp5.7.5IB Multiple Choice - Reaction Quotient & Kp5.7.6Extended Response - Kp & Bond Enthalpy
6

Reactivity - The Mechanisms of Chemical Change

6.1

Proton Transfer Reactions

6.1.1Brønsted-Lowry Theory
6.2

The pH Scale

6.2.1pH & pOH6.2.2Ionic Product of Water6.2.3Using Kw6.2.4IB Multiple Choice - pH & Kw6.2.5Extended Response - Kw
6.3

Strong & Weak Acids and Bases

6.3.1Strong Acids & Bases6.3.2Weak Acids & Ka6.3.3Weak Bases & Kb6.3.4Weak Acids & Bases Calculation Walkthrough
6.4

Acid Deposition

6.4.1Acid Deposition
6.5

Types of Organic Reactions

6.5.1Electrophilic Substitution
6.6

Oxidation & Reduction

6.6.1Oxidation Number Method6.6.2Winkler Method6.6.3Redox Reactions6.6.4Oxidation States & Equations6.6.5End of Topic Quiz - Redox6.6.6Extended Response - Redox Reactions
6.7

Electrochemical Cells

6.7.1Introduction to Electrochemical Cells6.7.2Galvanic Cells6.7.3Cell Notation6.7.4Electrolytic Cells6.7.5Extended Response - Voltaic Cells & Redox
6.8

Titration

6.8.1Redox Titrations6.8.2Titration Calculations6.8.3Atom Economy
6.9

Acid-Base Titrations

6.9.1Titration Calculation Weak Acid & Strong Base6.9.2Titration Experimental Detail6.9.3Extended Response - Titration6.9.4Titration Calculations6.9.5Titration Curves6.9.6Titration Calculation Strong Acid & Weak Base6.9.7IB Multiple Choice - Titrations6.9.8Polyprotic Acids6.9.9Titration Calculations Strong Acid & Strong Base6.9.10Titrations Curves 2
7

Measurement, Data Processing & Analysis

7.1

Uncertainties & Errors in Measurements & Results

7.1.1Qualitative & Quantitative data
7.2

Graphical Techniques

7.2.1Drawing & Interpreting Graphs7.2.2Linear Interpolation
7.3

Spectroscopic Identification of Organic Compounds

7.3.1Understanding NMR7.3.2Carbon-13 NMR7.3.3Hydrogen-1 NMR7.3.4Hydrogen-1 NMR 27.3.5End of Topic Quiz - NMR
7.4

Infrared Spectroscpy

7.4.1Infrared Spectroscopy

Practice questions on Balancing Equations & Ionic Equations

Can you answer these? Test yourself with free interactive practice on Seneca — used by over 10 million students.

  1. 1
    What volume of hydrogen gas will react with 2 liters of chlorine gas given that the reactant ratio is 1:1?Multiple choice
  2. 2
    What state symbol is correct for each of these substances at room temperature?True / false
  3. 3
    What do we call the letters that tell us the state of matter of substances in a chemical equation?Multiple choice
  4. 4
    Which of the following is the best definition of "net ionic equation"? Multiple choice
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IB Multiple Choice - Solids, Liquids, & Gases

Empirical & Molecular Formulas