Atomic Structure
Isotopes
Wild Card
Periodic Table Trends
Introduction to Bonding
Compounds and Mixtures
Bonding and Properties
Chemical Reactions
Exothermic and Endothermic
Mole Concept
100
What are the three subatomic particles in an atom?
Proton, neutron, and electron.
100
What is an isotope?
Atoms of the same element with different numbers of neutrons.
100
What is the symbol for a gas in a chemical equation?
(g)
100
What does each row on the periodic table represent?
A period — indicating the number of electron shells.
100
What is a chemical bond?
An electrostatic attraction between atoms or ions.
100
What’s the difference between an element and a compound?
An element is one type of atom; a compound is two or more chemically bonded.
100
What type of bonding occurs in metals?
Metallic Bonding
100
What is a chemical reaction?
A process that changes reactants into new products.
100
What is an exothermic reaction?
A reaction that releases heat.
100
How many particles are in one mole?
6.022 × 10²³
200
What part of the atom contains nearly all of its mass?
The nucleus.
200
Do isotopes have the same chemical or physical properties?
Same chemical, different physical properties.
200
Which element has the electron configuration [Ne]3s²3p⁴?
Sulphur (S)
200
What happens to atomic radius across a period and why?
It decreases due to increasing nuclear charge.
200
What type of bond forms between a metal and a non-metal?
Ionic bond.
200
What is a mixture?
A physical combination of substances that retain their properties.
200
Why do ionic compounds have high melting points?
Strong electrostatic forces between ions.
200
Name the products of a combustion reaction.
Carbon dioxide and water.
200
What happens to temperature in an endothermic reaction?
It decreases (heat is absorbed).
200
What’s the formula to calculate number of moles?
n = m / M
300
What does the atomic number of an element represent?
The number of protons.
300
How are isotopes written in nuclear notation?
300
What charge does a group 2 metal form?
+2
300
What is electronegativity?
An atom’s ability to attract shared electrons in a bond.
300
What does the term “valency” refer to?
The number of bonds an atom can form.
300
Give an example of a homogeneous mixture.
Saltwater
300
Are covalent molecules good conductors? Why or why not?
No, because they do not have free-moving electrons.
300
Write a balanced equation for: HCl + NaOH →
HCl + NaOH → NaCl + H₂O
300
Sketch an enthalpy diagram for an exothermic reaction.
Sketch an enthalpy diagram for an exothermic reaction.
300
What is the molar mass of H₂O?
18.02 g/mol
400
Write the electron configuration of sodium (Na).
1s² 2s² 2p⁶ 3s¹
400
Why do isotopes of an element behave similarly in chemical reactions?
They have the same electron configuration.
400
Why do noble gases not form compounds easily?
They have a full valence shell and are stable.
400
Which group contains the alkali metals?
Group 1
400
Define an ion.
An atom or group of atoms with a charge due to lost or gained electrons.
400
What does the boiling point of a mixture depend on?
The identity and relative amounts of substances in it.
400
Describe the structure of diamond
A 3D giant covalent lattice with each carbon bonded to 4 others.
400
What type of reaction is: Cu + AgNO₃ → Cu(NO₃)₂ + Ag?
Single displacement (redox).
400
What is bond enthalpy?
The energy needed to break one mole of a bond in the gas phase.
400
Calculate the number of moles in 10 g of NaCl. (M = 58.5 g/mol)
n = 10 ÷ 58.5 = 0.171 mol
500
Why does the first ionisation energy increase across a period?
Increasing nuclear charge pulls electrons closer, requiring more energy to remove one.
500
What is the relative atomic mass of an element based on?
The weighted average of its isotopes relative to 1/12 of a carbon-12 atom.
500
Write the empirical formula for a compound that is 75% carbon and 25% hydrogen by mass.
CH₄
500
Why do halogens become less reactive down the group?
Atomic radius increases, so attraction to electrons weakens.
500
Write the IUPAC name for Na₂SO₄.
Sodium sulfate.
500
Identify a method to separate sand and saltwater.
Filtration followed by evaporation.
500
Why are metals malleable?
Layers of atoms can slide over each other due to delocalised electrons.
500
What do state symbols (s), (l), (g), and (aq) mean?
Solid, liquid, gas, and aqueous (dissolved in water).
500
Explain why some reactions are exothermic using bond energies.
More energy is released forming bonds than used breaking them.
500
Define limiting reagent.
The reactant that is fully consumed and limits the amount of product formed.
600
Using the Aufbau principle, Hund's rule, and the Pauli exclusion principle, determine the full electron configuration for Ca²⁺.
1s² 2s² 2p⁶ 3s² 3p⁶
600
Chlorine has two isotopes: Cl-35 (75%) and Cl-37 (25%). Calculate its relative atomic mass.
(35×0.75)+(37×0.25)=35.5
600
Which category do enthalpy, ΔH fall under?
Exothermic and Endothermic Reactions
600
Compare the acidity of period 3 oxides from left to right.
They change from basic to amphoteric to acidic.
600
Draw the Lewis structure for H₂O and identify lone/bonding pairs.
2 bonding pairs (H–O–H), 2 lone pairs on oxygen.
600
Differentiate between a pure substance and a mixture using melting point data.
Pure substances melt at a sharp, fixed temperature; mixtures do not.
600
Describe the difference between covalent molecules and covalent networks.
Molecules are discrete entities; networks are extended lattices.
600
Balance this double displacement (precipitation) reaction and identify the precipitate:
BaCl2(aq)+Na2SO4(aq)→
BaCl2(aq)+Na2SO4(aq)→BaSO4(s)+2NaCl(aq). Precipitate = Barium Sulphate.
600
A 100.0 g sample of water is heated from 25.0°C to 60.0°C. Calculate the energy absorbed.
(Specific heat capacity of water = 4.18 J/g°C)
Q = mcΔT
= 100.0 × 4.18 × (60.0 – 25.0)
= 100.0 × 4.18 × 35.0
= 14,630 J or 14.63 kJ
600
Calculate % yield if theoretical = 5.0g and actual = 4.2g.
(4.2 ÷ 5.0) × 100 = 84%