Chemical Reactions: Test Review

Key facts

Where marks get lost

• ×Reading the Roman numeral as the number of metal atoms: iron(III) oxide is ${\text{Fe}}_2{\text{O}}_3$, not ${\text{Fe}}_3\text{O}$. The numeral is the charge on each iron ion. (Naming and Writing Formulas for Ionic Compounds)
• ×Forgetting to reduce: magnesium oxide is MgO, not ${\text{Mg}}_2{\text{O}}_2$. Ionic formulas always show the lowest ratio. (Naming and Writing Formulas for Ionic Compounds)
• ×Keeping the element name for the anion, as in "sodium chlorine". The non-metal ion takes the -ide ending: sodium chloride. (Naming and Writing Formulas for Ionic Compounds)
• ×Adding Roman numerals to metals with only one charge, such as "sodium(I) chloride" or "zinc(II) oxide". Group 1, group 2, aluminum, zinc and silver never need them. (Naming and Writing Formulas for Ionic Compounds)
• ×Writing the charges into the finished formula, like ${\text{Ca}}^{2+}{\text{Cl}}_2^{-}$. Charges are used to work out the formula, then dropped. (Naming and Writing Formulas for Ionic Compounds)
• ×Leaving out brackets: calcium hydroxide is ${\text{Ca(OH)}}_2$, not ${\text{CaOH}}_2$. Without brackets the 2 applies only to the H. (Polyatomic Ions, Molecular Compounds and Acids)
• ×Changing a polyatomic ion to balance charges, such as writing $\text{NO}$ or ${\text{NO}}_6$ instead of using two ${\text{NO}}_3$ groups. The ion's own formula never changes. (Polyatomic Ions, Molecular Compounds and Acids)
• ×Using prefixes for ionic compounds: ${\text{CaCl}}_2$ is calcium chloride, not calcium dichloride. Prefixes are only for two non-metals. (Polyatomic Ions, Molecular Compounds and Acids)
• ×Putting mono on the first element ("monocarbon dioxide") or forgetting it on the second ("carbon oxide" for CO). (Polyatomic Ions, Molecular Compounds and Acids)
• ×Calling ${\text{H}}_2{\text{SO}}_4$ "hydrosulfuric acid". The hydro- prefix is only for binary acids with no oxygen; an -ate ion gives an -ic acid with no hydro-. (Polyatomic Ions, Molecular Compounds and Acids)
• ×Treating every bubble as proof of a reaction. Boiling water bubbles too; bubbles count only when a new gas forms. (Evidence of Chemical Change and Conservation of Mass)
• ×Saying "mass was lost" when a gas escapes from an open container. The mass left the system; it was not destroyed. (Evidence of Chemical Change and Conservation of Mass)
• ×Forgetting that a reactant can come from the air: burning steel wool gains mass because oxygen joins the iron. (Evidence of Chemical Change and Conservation of Mass)
• ×Mixing up reactants and products: reactants are on the left of the arrow (what you start with), products on the right (what forms). (Evidence of Chemical Change and Conservation of Mass)
• ×Using a single clue as proof. A colour change alone, like adding food colouring, can be physical; look for several kinds of evidence. (Evidence of Chemical Change and Conservation of Mass)
• ×Changing a subscript to balance, such as turning ${\text{H}}_2\text{O}$ into ${\text{H}}_2{\text{O}}_2$. That makes a different substance; only coefficients may change. (Word Equations and Balanced Chemical Equations)
• ×Writing oxygen, hydrogen or chlorine as single atoms (O, H, Cl) when they appear as elements. On their own they are diatomic: ${\text{O}}_2$, ${\text{H}}_2$, ${\text{Cl}}_2$. (Word Equations and Balanced Chemical Equations)
• ×Forgetting that a coefficient multiplies the whole formula: $2{\text{H}}_2\text{O}$ has 4 H and 2 O, not 4 H and 1 O. (Word Equations and Balanced Chemical Equations)
• ×Not recounting every element after each change, so balancing one element quietly unbalances another. (Word Equations and Balanced Chemical Equations)
• ×Leaving coefficients that can be reduced, such as 4, 2, 4 instead of 2, 1, 2. (Word Equations and Balanced Chemical Equations)
• ×Classifying by how complicated the formulas look instead of counting: one product means synthesis, one reactant means decomposition. (Synthesis and Decomposition Reactions)
• ×Writing the product of a metal and a non-metal without using ion charges, such as "AlCl" instead of ${\text{AlCl}}_3$. (Synthesis and Decomposition Reactions)
• ×Forgetting that oxygen, hydrogen, nitrogen and the halogens are diatomic when they are products of a decomposition (writing O instead of ${\text{O}}_2$). (Synthesis and Decomposition Reactions)
• ×Mixing up which oxide gives which: metal oxides with water give bases, and non-metal oxides with water give acids. (Synthesis and Decomposition Reactions)
• ×Predicting a single displacement without checking the activity series. If the free metal is below the metal in the compound, the answer is "no reaction". (Single Displacement, Double Displacement and Combustion)
• ×Copying old subscripts into double displacement products, such as writing ${\text{Na}}_2\text{Cl}$ instead of NaCl. Rebuild each new formula from the ion charges. (Single Displacement, Double Displacement and Combustion)
• ×Swapping the wrong parts in a double displacement, like pairing two cations. A cation always pairs with an anion. (Single Displacement, Double Displacement and Combustion)
• ×Forgetting that oxygen is ${\text{O}}_2$ in combustion, or forgetting water as a product when the fuel contains hydrogen. (Single Displacement, Double Displacement and Combustion)
• ×Balancing oxygen before carbon and hydrogen in a combustion equation; doing oxygen last saves a lot of rework. (Single Displacement, Double Displacement and Combustion)
• ×Thinking a higher pH means more acidic. It is the reverse: the lower the pH, the more acidic. (Acids, Bases, pH and Indicators)
• ×Treating the pH scale as linear, e.g. saying pH 2 is "twice as acidic" as pH 4. A difference of 2 units is $10^2=100$ times. (Acids, Bases, pH and Indicators)
• ×Mixing up litmus colours. Remember "acid turns blue litmus red"; base turns red litmus blue. (Acids, Bases, pH and Indicators)
• ×Concluding that a solution is acidic because phenolphthalein stayed colourless. It also stays colourless in neutral and slightly basic solutions (below about pH 8.2). (Acids, Bases, pH and Indicators)
• ×Identifying acids or bases by taste, touch or smell. Use indicators or a pH meter. (Acids, Bases, pH and Indicators)
• ×Thinking "salt" always means sodium chloride. In chemistry a salt is any ionic compound made from the cation of a base and the anion of an acid, such as ${\text{K}}_2{\text{SO}}_4$ or ${\text{CaCl}}_2$. (Neutralization and Reactions in the Environment)
• ×Building the salt from the wrong ions, such as combining the acid's hydrogen with the base's hydroxide into the salt. H⁺ and OH⁻ become water; the other two ions form the salt. (Neutralization and Reactions in the Environment)
• ×Forgetting to balance the water: an acid with two H (like ${\text{H}}_2{\text{SO}}_4$) needs two OH⁻ and makes two water molecules. (Neutralization and Reactions in the Environment)
• ×Assuming a neutralized solution always has a pH of exactly 7. Exactly matched amounts of a strong acid and a strong base (like HCl and NaOH) give about pH 7, but with a weak acid or weak base (like vinegar or ammonia) the salt solution can end up a little above or below 7, and an excess of either reactant shifts the pH further. (Neutralization and Reactions in the Environment)
• ×Treating an acid spill with a strong base. A weak base such as baking soda or limestone is safer because an excess does not create a new hazard. (Neutralization and Reactions in the Environment)

Practice test

24 questions from across the unit. Each one shows the lesson it comes from.

Endless drill

Fresh numbers every time, for the calculation questions in this unit.