QCAA Chemistry 2023 Units 1-4
🇬🇧
In English
In English
Practice Known Questions
Stay up to date with your due questions
Complete 5 questions to enable practice
Exams
Exam: Test your skills
Test your skills in exam mode
Learn New Questions
Manual Mode [BETA]
Select your own question and answer types
Specific modes
Learn with flashcards
Complete the sentence
Listening & SpellingSpelling: Type what you hear
multiple choiceMultiple choice mode
SpeakingAnswer with voice
Speaking & ListeningPractice pronunciation
TypingTyping only mode
QCAA Chemistry 2023 Units 1-4 - Leaderboard
QCAA Chemistry 2023 Units 1-4 - Details
Levels:
Questions:
171 questions
| 🇬🇧 | 🇬🇧 |
| Define electronegativity | The tendency for an atom of a given chemical element to attract shared electrons when forming a chemical bond |
| Define ionisation energy | The amount of energy required to remove an electron from an isolated atom or molecule. |
| Define valency | The ability of an atom to form chemical bonds with other atoms |
| What is the trend of ionisation energy? | Ionization energy increases moving from left to right across an element period (This is due to valence shell stability). Ionization energy of the elements within a group generally decreases from top to bottom (This is due to electron shielding). |
| What is the trend of electronegativity | Electronegativity decreases down a group and increases across a period |
| What is the trend of valency? | Valency first increases and then decreases as we go from left to right in a period but remains the same in a group. |
| Define a metallic element | Any chemical elements that are usually shiny solids that conduct heat or electricity and can be formed into sheets etc |
| Define non-metallic elements | A chemical element (such as boron, carbon, or nitrogen) that lacks the characteristics of a metal. |
| Define an alkali metal | Very reactive chemical species that readily lose their one valence electron to form ionic compounds with nonmetal elements |
| Define a halogen | They are reactive non-metallic elements that form strongly acidic compounds with hydrogen. |
| Define Aufbau's principle | An electron enters the orbital with lowest energy first and subsequent electrons are fed in the order of increasing energies. |
| Define Hund's rule | Every orbital is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin. |
| Define Pauli's exclusion principle | No more than two electrons can occupy the same orbital and two electrons in the same orbital must have opposite spins |
| Define electrostatic forces | The attractive or repulsive force between two electrically charged objects. |
| Define valence electrons | The electrons in the outermost shell, or energy level, of an atom. |
| Define valency | The number of other atoms with which an atom of the element can combine. |
| Define electron lone pairs | A pair of valence electrons that in a covalent bond are not exchanged with another atom |
| What are the types of bonding? | Ionic, covalent and metallic |
| Define ionic bonding | A chemical connection in which one atom loses valence electrons and gains them from another |
| Define covalent bonding | The mutual sharing of one or more pairs of electrons between two atoms. |
| Define metallic bonding | A chemical bond formed between positively charged atoms in which the free electrons are shared among a lattice of cations. |
| What is the formula for relative atomic mass (RAM)? | ∑ isotope mass x isotope abundance / 100. |
| What is the isotope abundance formula? | Average atomic mass/∑ isotope mass |
| Define a pure substance | A form of matter that has a constant composition and properties that are constant throughout the sample. |
| Define a heterogenous mixture | A mixture in which the composition is not uniform throughout the mixture. |
| Define a homogenous mixture | A mixture in which the composition is uniform throughout the mixture. |
| Define distillation | The process involving the conversion of a liquid into vapour that is subsequently condensed back to liquid form |
| Define an ionic compound | Neutral compounds made up of positively charged ions called cations and negatively charged ions called anions |
| Define an allotrope | The existence of a substance and especially an element in two or more different forms usually in the same phase. |
| Define a metallic lattice | The type of bond that is formed to create the structure of metals |
| Define giant covalent networks | A three-dimensional structure of atoms that are joined by covalent bonds. Substances with giant covalent structures are solids with very high melting points |
| Define a hydrocarbon | An organic compound containing only carbon and hydrogen. |
| Define an alkane | Alkanes are saturated hydrocarbons. This means that their carbon atoms are joined to each other by single bonds. This makes them relatively unreactive |
| Define a alkene | Alkenes are a homologous series of hydrocarbons that contain a carbon-carbon double bond. |
| Define a single displacement reaction | A reaction in which one element is substituted for another element in a compound |
| Define a double displacement reaction | Where two compounds react, and positive ions and the negative ions of the two reactants switch places, forming two new compounds or products. |
| Define a combustion reaction | A reaction in which a substance reacts with oxygen gas, releasing energy in the form of light and heat |
| Define a decomposition reaction | Processes in which chemical species break up into simpler parts |
| Define a redox reaction | A chemical reaction that takes place between an oxidizing substance and a reducing substance. |
| Define a combination reaction | A reaction in which two reactants combine to form one product |
| State the law of conservation of energy | The amount of energy is neither created nor destroyed. |
| Define kinetic energy | The energy of motion, observable as the movement of an object, particle, or set of particles |
| Define an exothermic reaction | When energy is transferred to the surroundings |
| Define an endothermic reaction | Chemical reactions in which the reactants absorb heat energy from the surroundings to form products. |
| Define bond enthalpies | The amount of energy stored in a bond between atoms in a molecule. |
| Define specific heat capacity | The energy required to increase temperature of material of a certain mass by 1°C |
| Define heat change | The transfer of energy from one body to another as a result of a difference in temperature or a change in phase. |
| State Avogadro's number | 6.02 × 1023 |
| Define molar mass | The sum of the total mass in grams of the atoms present to make up a molecule per mole |
| Define a mole | The amount of substance containing the same number of molecules as there are atoms in exactly 12 g of 12C. |
| Define relative atomic mass | The weight in grams of the number of atoms of the element contained in 12.00 g of carbon-12 |
| Define empirical formula | Chemical formula showing the simplest ratio of elements in a compound rather than the total number of atoms in the molecule |
| Define molecular formula | Chemical formula that gives the total number of atoms of each element in each molecule of a substance |
| Define the VSEPR theory | A model used to predict 3-D molecular geometry based on the number of valence shell electron bond pairs among the atoms in a molecule or ion |
| What are the primary molecular shapes? | Linear, bent, trigonal planar, tetrahedral, pyramidal |
| Define dispersion forces | A force acting between atoms and molecules that are normally electrically symmetric and are the weakest intermolecular attractive forces. |
| Define dipole-dipole attraction | Attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. |
| Define hydrogen bonding | An attraction between a hydrogen atom in one polar molecule and a small electronegative atom in usually another molecule of the same or a different polar substance. |
| State STP (standard temperature and pressure) | Temperature of 273K (0 °C) and an absolute pressure of exactly 105 Pa |
| Define the kinetic theory of gases | That all matter is made of small particles that are in random motion and that have space between them. |
| State the ideal gas law | PV = nRT (P = pressure, V= volume, n= amount of substance, R=ideal gas constant, T=temperature) |
| Define Boyle's law | That the pressure of a given mass of an ideal gas is inversely proportional to its volume at a constant temperature. |
| Define Charles law | That the volume of an ideal gas at constant pressure is directly proportional to the absolute temperature. |
| Define Gay Lussac's law | That the volumes of gases undergoing a reaction at constant pressure and temperature are in a simple ratio to each other and to that of the product. |
| Define Avogadro's law | That equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. |
| Define the relation between pressure and temperature | The pressure of a given amount of gas is directly proportional to the temperature at a given volume. |
| Define the relation between pressure and volume | The volume of a fixed amount of a gas is inversely proportional to its pressure. |
| Define the relation between temperature and volume | The volume of a given amount of gas is directly proportional to its temperature |
| Define molarity | The moles of a solute per liters of a solution |
| Define a solution | A type of homogeneous mixture composed of two or more substances. |
| Define a solvent | A substance that dissolves a solute, resulting in a solution. |
| Define a saturated solution | A solution in which the maximum amount of solvent has been dissolved |
| Define an unsaturated solution | A solution in which more solute can be dissolved. |
| Define a supersaturated solution | When the concentration of a solute exceeds the concentration specified by the value of solubility at equilibrium. |
| Define a precipitate | A substance separated from a solution or suspension by chemical or physical change |
| Define solubility | The ability to be dissolved, especially in water. |
| Define a solubility curve | Graphic representation of the variation with changing temperature of the solubility of a given substance in a given solvent. |
| Define pH | A quantitative measure of the acidity or basicity of aqueous or other liquid solutions |
| Define an Arrhenius base | Any species that increases the concentration of hydroxide ions |
| Define an Arrhenius acid | A compound that increases the Hydronium ion concentration in aqueous solution. |
| Define collision theory | Theory used to predict the rates of chemical reactions, particularly for gases |
| Define Maxwell-Boltzmann distribution | Describes the distribution of speeds among the particles in a sample of gas at a given temperature. |
| Define activation energy | The minimum quantity of energy which the reacting species must possess in order to undergo a specified reaction. |
| Define energy profile diagrams | Shows whether a reaction is exothermic or endothermic |
| Define an open chemical system. | Allowing matter and energy to be exchanged with the surroundings. |
| Define a closed chemical system. | Allowing energy, but not matter, to be exchanged with the surroundings. |
| Define chemical equilibrium. | The state in a reversible chemical reaction where the rate of the forward reaction is equal to the reverse reaction, resulting in a stable concentration of reactants and products over time. |
| What factors affect equilibrium? | Temperature, pressure, concentrations of reactants and products, and the presence of catalysts. |
| What is the effect of temperature change on chemical systems at equilibrium? | Considering the enthalpy change for the forward and reverse reactions; an endothermic forward reaction will shift to the right with increasing temperature, while an exothermic forward reaction will shift to the left. |
| Define Le Châtelier’s principle. | If a system in equilibrium is subjected to a change in temperature, pressure, or concentration, the system will adjust to counteract the change and restore equilibrium. |
| What is the effect of changing concentration and pressure on chemical systems at equilibrium (applying collision theory). | Shifts the chemical equilibrium of a reaction by affecting the rate of forward and reverse reactions based on collision theory; increasing concentration or pressure favors side with less moles of gas, while decreasing concentration or pressure shifts the equilibrium in the direction to oppose the change. |
| State the formula for equilibrium constant (Kc) | ?c =[C]^c x [D]^d / [A]^a x [B]^b |
| Deduce the extent of a reaction from the magnitude of the equilibrium constant. | Larger values indicate a more complete conversion of reactants to products at equilibrium. |
| Define an acid (according to the Brønsted–Lowry definition) | Acids are substances that act as proton donors in a chemical reaction. |
| Define a base (according to the Brønsted–Lowry definition) | Bases are substances that act as proton receivers in a chemical reaction. |
| Define the term monoprotic | "Monoprotic" refers to an acid or base that can donate or accept only one proton (H+) in a chemical reaction. |
| Define the term polyprotic | "Polyprotic" refers to an acid or base that can donate or accept more than one proton (H+) in a chemical reaction. |
| Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity. | Strong acids and bases dissociate completely in water, leading to a high degree of ionization, strong reaction with water, and high electrical conductivity. Weak acids and bases partially dissociate, resulting in lower ionization, weaker reaction with water, and lower electrical conductivity. |
| Define what a 'strong' acid and base is. | Strong acids and bases fully ionize or dissociate in water, resulting in high concentrations of hydrogen ions (H+) for strong acids and hydroxide ions (OH-) for strong bases |
| Define what is meant by 'concentrated acids and bases'. | Concentrated acids and bases refer to solutions with a large amount of acid or base dissolved in water regardless of their ionization or dissociation properties. |