All of AQA CHEMISTRY Paper 1 in 30 minutes - GCSE Science Revision

This comprehensive overview of AQA GCSE Chemistry Paper 1 covers essential topics including atoms, bonding, quantitative chemistry, and chemical and energy changes, applicable to both higher and foundation tiers. It explains the structure and properties of atoms, the periodic table, ionic and covalent bonding, and the behavior of metals and non-metals. The guide also delves into chemical reactions, balancing equations, states of matter, separation techniques, and the history of atomic models. Additionally, it touches on quantitative chemistry concepts like moles and yields, the reactivity series, electrolysis, and energy changes in reactions, providing a detailed yet succinct resource for exam preparation.

• Substances are made of atoms, which are represented in the periodic table by symbols.
• A compound consists of two or more different types of atoms chemically bonded together.
• Chemical formulas represent compounds, e.g., H₂O for water, indicating two hydrogen atoms and one oxygen atom.
• Atoms are not created or destroyed in chemical reactions; equations must be balanced to reflect this.

"A compound is a substance that contains two or more different types of atoms chemically bonded together."

• Explanation: Defines what a compound is in terms of chemical bonding.

"Atoms are not created or destroyed in any chemical reaction there must be the same number of each type of atom on both sides so sometimes we must balance equations."

• Explanation: Emphasizes the principle of conservation of mass in chemical reactions and the necessity of balancing equations.

• Chemical reactions can be represented by word equations and chemical equations using symbols.
• Balancing equations involves ensuring the same number of each type of atom on both sides of the equation.
• Start by balancing atoms that are only in compounds, then move to elements like oxygen.

"We can represent a reaction with a word equation and a chemical equation using symbols as atoms are not created or destroyed in any chemical reaction there must be the same number of each type of atom on both sides."

• Explanation: Highlights the importance of representing chemical reactions accurately and the need for balanced equations.

"Pro tip: start balancing atoms that are only in compounds so with this one let's go with the carbons first."

• Explanation: Provides a strategy for balancing chemical equations more efficiently.

• Mixtures are combinations of different elements and compounds that are not chemically bonded.
• Techniques to separate mixtures include filtration, crystallization, distillation, and fractional distillation.
• These processes are physical changes, not chemical reactions, as no new substances are formed.

"A mixture is any combination of any different types of elements and compounds that aren't chemically bonded together for example air is a mixture of oxygen nitrogen and more."

• Explanation: Defines what a mixture is and gives an example.

"Filtration can separate large insoluble particles from a liquid, crystallization can leave a solute behind after evaporating the solvent, and distillation involves heating and cooling to separate substances."

• Explanation: Describes different methods to separate components of mixtures.

• The three main states of matter are solid, liquid, and gas.
• Physical changes, such as melting and evaporation, involve changes in energy but do not form new substances.
• State symbols in chemical equations indicate the physical state of substances (s for solid, l for liquid, g for gas, aq for aqueous).

"Solid liquid and gas are the three main states of matter for example water can be ice a solid where the particles or molecules in this case vibrate around fixed positions."

• Explanation: Describes the states of matter using water as an example.

"To melt or evaporate a substance you must supply energy usually in the form of heat to overcome the electrostatic forces of attraction between the particles."

• Explanation: Explains the energy changes involved in physical state changes.

• Early models of the atom included the plum pudding model by JJ Thompson and the nuclear model by Ernest Rutherford.
• Niels Bohr discovered that electrons exist in shells, and James Chadwick discovered neutrons in the nucleus.
• Protons and neutrons have a relative mass of one, while electrons have a very small mass.

"JJ Thompson discovered that atoms are made up of positive and negative charges he came up with the plum pudding model of the atom a positive charge with lots of little electrons dotted around it."

• Explanation: Describes the plum pudding model of the atom.

"Ernest Rutherford found that the positive charge must actually be incredibly small We Now call this the nucleus and the electron must orbit relatively far away from it."

• Explanation: Explains Rutherford's discovery of the nucleus and the nuclear model of the atom.

• The periodic table provides information about elements, including atomic number (number of protons) and mass number (protons plus neutrons).
• Isotopes are atoms of the same element with different numbers of neutrons.
• The periodic table was organized by Dmitri Mendeleev based on element properties, leading to the modern periodic table.
• Electron shells fill from the inside out, with specific numbers of electrons in each shell.

"The periodic table tells us everything we need to know about an atom the bottom number is the atomic number that's the number of protons in the nucleus this is what determines what element you have."

• Explanation: Describes how the periodic table provides essential information about elements.

"Dimitri Mev then came along and grouped elements together based on their properties even if the order didn't follow atomic weight using this method he found there were gaps in his table he asserted that these elements were yet to be discovered."

• Explanation: Explains Mendeleev's contribution to the periodic table and his predictions about undiscovered elements.

• Metals are found to the left of the staircase on the periodic table and tend to donate electrons.
• Non-metals are to the right of the staircase and tend to accept electrons.
• Atoms that gain or lose electrons become ions, which have a net charge.

"Everything to the left of this staircase is called a metal metal atoms always donate electrons to gain an empty outer shell of electrons."

• Explanation: Describes the general behavior of metals in terms of electron donation.

"To the right of the staircase nonmetals they always accept electrons to gain a full outer shell."

• Explanation: Describes the general behavior of non-metals in terms of electron acceptance.

• The number of electrons in an atom's outer shell determines its chemical properties.
• Transition metals can donate different numbers of electrons when they bond to different substances.

"The transition metals work in a really weird way so they don't get their own group. In fact, it turns out this is because they can donate a different number of electrons when they bond to different things."

• Transition metals exhibit variable oxidation states due to their ability to donate different numbers of electrons.

• Alkali metals have one electron in their outer shell, which they donate when bonding.
• Reactivity increases down the group due to weaker electrostatic attraction between the outer electron and the nucleus.

"The atoms in group one are called the alkali metals. They all have one electron in their outer shell which they give away when they bond to something, so they have similar properties like when they react with water. The further down the group you go, the further that outer electron is from the nucleus, so the electrostatic attraction is weaker between the negative electron and the positive nucleus."

• Alkali metals become more reactive as you go down the group.

• Halogens have seven electrons in their outer shell and need one more to complete it.
• Reactivity decreases down the group as the outer shell is further from the nucleus.
• Boiling points increase down the group.

"Group seven are what we call the halogens. They have seven electrons in their outer shell, so they need one more to gain a full outer shell. The further down the group you go, the less readily an electron is accepted onto that shell that's further away from the nucleus, so they get less reactive down the group. Their boiling points also increase down the group too."

• Halogens become less reactive and have higher boiling points as you go down the group.

• Noble gases have a full outer shell and are very unreactive.
• Helium is included in this group despite having only two electrons.

"Group zero, sometimes referred to as group eight, are called the noble gases. They already have a full outer shell, so they don't react. In reality, they can react under special conditions, so we just say they're very unreactive."

• Noble gases are stable and rarely react due to their full outer shell.

• Metals lose electrons to form positive ions (cations).
• Non-metals gain electrons to form negative ions (anions).

"Metals become positively charged when they lose them; they always form positive ions. All of group one lose one electron when they turn into an ion, so all of their ions are one plus. Group two lose two electrons to get an empty outer shell, so their ions are all two plus. Group seven gain one electron each, so all their ions are minus. Group six's ions are all two minus."

• The charge of an ion depends on the number of electrons lost or gained to achieve a full outer shell.

• Transition metals can form multiple ions with different charges.
• Example: Iron can form Fe2+ or Fe3+.

"Transition metals can donate different numbers of electrons. For example, an iron ion can be Fe2+ or Fe3+; it can donate two or three electrons, so we give them the names iron 2 and iron 3 to distinguish between them."

• Transition metals have variable oxidation states, allowing them to form different ions.

• Metal atoms bond through metallic bonding, forming a lattice of ions surrounded by delocalized electrons.
• Metals are good conductors of electricity and heat due to the mobility of delocalized electrons.

"Metal atoms bond to each other through metallic bonding. Essentially, a lattice or grid of ions is formed with a sea of delocalized electrons around them. Delocalized just means they're not exactly on the atom. As these electrons are free to move, metals make good conductors of electricity and heat."

• Metallic bonding involves a lattice of ions and delocalized electrons, enabling conductivity.

• Metals bond to non-metals through ionic bonding, involving the transfer of electrons.
• Example: Lithium donates its outer electron to chlorine, forming Li+ and Cl- ions.

"Metals bond to non-metals through ionic bonding. For example, a lithium atom donates its outer electron to the chlorine. We can draw a dot and cross diagram to show where the electrons end up."

• Ionic bonding involves the transfer of electrons, resulting in the formation of cations and anions.

• Ionic compounds form a lattice structure and have high melting and boiling points.
• They conduct electricity when molten or dissolved in solution.

"Ionic compounds consist of lots of repeating units of these ions in a lattice to form a crystal. They have high melting points and boiling points due to the strong electrostatic forces that need to be overcome. They can conduct electricity but only in liquid form, that is molten or when dissolved in solution."

• Ionic compounds have high melting points and can conduct electricity in liquid form due to free-moving ions.

• Non-metals bond to each other through covalent bonding, sharing electrons to achieve full outer shells.
• Example: Chlorine gas (Cl2) involves each chlorine atom sharing an electron.

"Non-metals bond to each other with covalent bonding to form molecules. They do this by sharing electrons to gain full outer shells. For example, chlorine gas is Cl2; each chlorine atom shares an electron with the other."

• Covalent bonding involves the sharing of electrons between non-metals to form molecules.

• Simple molecular structures have low boiling points due to weak intermolecular forces.
• They do not conduct electricity, even as liquids.

"Simple molecular structures have relatively low boiling points as there are only weak intermolecular forces between them that need to be overcome with heating. Unlike ionic compounds, these can't conduct electricity even as liquids."

• Simple molecular structures have weak intermolecular forces and do not conduct electricity.

• Giant covalent structures form through extensive covalent bonding, creating one large molecule.
• Example: Diamond, a crystal of carbon atoms, is extremely hard with a high melting point.

"Giant covalent bonding is similar to the lattice nature of ionic compounds. Atoms form covalent bonds to other atoms, which form bonds to other atoms, and so on until what we have is one giant molecule. Diamond is an example of this; it's a crystal of carbon atoms bonded to each other."

• Giant covalent structures, like diamond, are extremely hard and have high melting points due to extensive covalent bonding.

• Graphite consists of layers of carbon atoms with three bonds each, allowing for electrical conductivity and easy sliding of layers.
• Allotropes of carbon include graphene and fullerenes.

"Graphite consists of layers of carbons with three bonds each in a hexagonal structure. The spare delocalized electrons form special weak bonds between the layers, which means that it can conduct electricity because the electrons can move between the layers as well. Graphene is just a single layer of graphite. Fullerenes are 3D structures of carbon atoms."

• Graphite conducts electricity and has layers that slide easily, while graphene and fullerenes are other forms of carbon with unique properties.

• The total mass of all substances is conserved in a chemical reaction.

"The total mass of all substances is conserved in a chemical reaction."

• Mass conservation is a fundamental principle in chemical reactions.

• Balancing equations ensures the conservation of mass in chemical reactions.
• Relative atomic mass (RAM) can be summed to find the relative formula mass of compounds.
• Example: CO2 has a relative formula mass of 44 (12 + 2x16).

"Some reactions produce a gas product which if it leaves the reaction vessel will result in a seeming decrease in mass of the reactants."

• Gas products leaving the reaction vessel can cause an apparent decrease in mass.

• A mole is a specific number of atoms or molecules, used to compare amounts of substances.
• One mole of a substance has a mass equal to its relative atomic or formula mass.
• Equation for moles: moles = grams / RAM.

"If you have as many grams of a substance as its relative atomic or formula mass, you have one mole."

• One mole of a substance equals its relative atomic or formula mass in grams.

• To balance methane combustion, use the ratio of 2 oxygen molecules per 1 methane molecule.
• Process to find mass of water produced from 64g of methane: Convert mass to moles, use stoichiometry, convert moles back to mass.

"Mass moles moles Mass we switch from one to the other at the halfway mark."

• Use mass-moles-moles-mass steps to find the mass of the product from a given reactant mass.

• Limiting reactants determine the extent of the reaction; the reactant that runs out first limits the reaction.
• Example: If oxygen is limited in methane combustion, not all methane will react.

"We say that the oxygen is the limiting reactant in this case; it ran out first."

• The limiting reactant is the one that is completely consumed first in the reaction.

• Concentration can be given in grams per decimeter cubed (g/dm³) or moles per decimeter cubed (mol/dm³).
• 1 mole of HCl in 1 dm³ of water has a concentration of 1 mol/dm³.

"Sometimes we shorten this to just one Moler."

• Concentration can be expressed as molarity (M), e.g., 1 Molar (1M).

• Reversible reactions like the Haber process reach equilibrium with some reactants left unreacted.
• Percentage yield = (actual yield / theoretical yield) x 100.

"Percentage yield merely tells you how much product is actually made compared to how much you could have made in theory had all the reactants reacted."

• Percentage yield indicates the efficiency of a reaction in producing the desired product.

• Atom economy measures the efficiency of mass usage in a reaction.
• Calculation: (RAM of desired product / total RAM of reactants) x 100.

"Atom economy tells you how much of a desired product you get out of a reaction compared to the mass of the reactants that went in."

• High atom economy is desirable for efficient and sustainable chemical processes.

• One mole of any gas occupies 24 dm³ at room temperature and pressure (RTP: 20°C and 1 atmosphere).

"One mole of any gas takes up a volume of 24 decim cubed regardless of its relative mass."

• Molar volume of gases at RTP is 24 dm³.

• Metals vary in reactivity; more reactive metals displace less reactive metals from compounds.
• Example: Zinc displaces copper from copper sulfate solution.

"A more reactive metal will displace a less reactive metal from a compound."

• Reactivity series helps predict outcomes of displacement reactions.

• Metals less reactive than carbon can be extracted from their ores by displacement with carbon (smelting).
• Reduction is the loss of oxygen or gain of electrons.

"Ion can be displaced from ion oxide with carbon; this is called smelting."

• Smelting involves reduction of metal oxides using carbon.

• Redox reactions involve oxidation (loss of electrons) and reduction (gain of electrons).
• Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

"Reduction and oxidation happen depending on whether a reactant loses or gains electrons."

• Redox reactions are fundamental in understanding electron transfer in reactions.

• Acids react with bases to form salts and water, neutralizing each other.
• Example: Sodium hydroxide reacts with hydrochloric acid to form sodium chloride and water.

"Sodium hydroxide and hydrochloric acid makes sodium chloride in water."

• Neutralization reactions result in the formation of a salt and water.

• Dissolved substances dissociate into ions; the pH scale measures the concentration of H+ ions.
• The pH scale is logarithmic, meaning each unit change represents a tenfold change in H+ concentration.

"An acid that has a pH of three will have 10 times the concentration of these compared to an acid with a pH of four."

• The pH scale indicates the acidity or basicity of a solution based on H+ ion concentration.

• pH scale measures the concentration of H+ ions in a solution.
• A lower pH indicates a higher concentration of H+ ions.
• Acids with pH3 have 100 times the concentration of H+ ions compared to acids with pH5.
• Alkaline substances work similarly but with OH- ions.

"An acid of pH4 pH3 would have a 100 times the concentration of H+ ions compared to an acid of pH5 and so on."

• The concentration of H+ ions increases exponentially as pH decreases.

• Strong acids dissociate completely in solution (e.g., Hydrochloric, Nitric, and Sulfuric acids).
• Weak acids only partially dissociate (e.g., Ethanoic, Citric, and Carbonic acids).
• pH of an acid depends on both its strength and concentration.

"A strong acid is one that dissociates or ionizes completely when in solution like Hydrochloric, Nitric, and Sulfuric acids."

• Strong acids fully ionize in solution, leading to a lower pH.

"Weak acids on the other hand only partially dissociate like Ethanoic, Citric, and Carbonic acids."

• Weak acids do not fully ionize, resulting in a higher pH compared to strong acids of the same concentration.

• Titrations are used to deduce the concentration of an acid or an alkaline.
• A known volume of alkali is measured using a glass pipet and placed in a conical flask with an indicator.
• Acid of unknown concentration is added from a burette until neutralization is indicated by a color change.

"We use a glass pipet to measure out a known volume of alkali and put it in a conical flask with a few drops of an indicator like methyl orange."

• The known volume of alkali and an indicator are crucial for the titration process.

• Electrolysis involves passing a current through a melted ionic compound or solution to cause a chemical change.
• Positive metal ions (cations) move to the negatively charged electrode (cathode) and are reduced.
• Negative ions (anions) move to the positively charged electrode (anode) and are oxidized.

"Cations are always reduced at the cathode so in this case solid aluminium is formed on the cathode."

• Reduction of cations at the cathode results in the formation of solid metal.

"Anions are always oxidized at the anode this is one way of purifying metals or extracting them from compounds."

• Oxidation of anions at the anode is used to purify or extract metals.

• Chemical reactions involve energy transfers: energy is needed to break bonds, and energy is released when bonds form.
• Exothermic reactions release more energy than they absorb, resulting in a temperature increase.
• Endothermic reactions absorb more energy than they release, resulting in a temperature decrease.

"If there is more energy released from bonds made than energy needed to break bonds, we say this is a net energy released and we should observe an increase in temperature as a result."

• Exothermic reactions result in a net release of energy and an increase in temperature.

"If it's the other way around, there is net energy input into the reaction so the reaction should get colder this is an endothermic reaction."

• Endothermic reactions result in a net absorption of energy and a decrease in temperature.

• Neutralization reactions can be used to measure energy changes.
• Insulated containers and thermometers are used to measure temperature changes during reactions.
• Energy profiles help visualize the energy differences between reactants and products.

"We carry out a neutralization reaction between an acid and alkaline in a polystyrene cup which is well insulated and a thermometer pokes through a lid that sits on top."

• Practical setups for measuring energy changes include insulated containers and thermometers.

• Ionic substances in solution can undergo electrolysis.
• Reactivity determines which ions are reduced or oxidized at the electrodes.
• Hydrogen gas is produced at the cathode if H+ ions are less reactive than metal ions.

"The more reactive ion stays in solution while the less reactive one moves to the electrode."

• Reactivity influences which ions are reduced or oxidized during electrolysis.

"If the metal is less reactive than hydrogen, say copper in copper sulfate solution, it forms on the cathode instead and the H+ ions stay in solution."

• Less reactive metals like copper will be reduced at the cathode instead of hydrogen.

• Energy profiles show the potential energy changes during reactions.
• Exothermic reactions have products with lower potential energy than reactants.
• Endothermic reactions have products with higher potential energy than reactants.

"The y-axis is potential energy and you should know that usually in science potential energy and kinetic energy do a balancing act if one goes down the other one goes up."

• Energy profiles illustrate the relationship between potential and kinetic energy in reactions.

"If the potential energy of the products is less than the reactants, they must have gained kinetic energy and that always means a hotter temperature this is an exothermic reaction."

• Exothermic reactions result in a decrease in potential energy and an increase in temperature.

• Cells and batteries produce voltage through chemical reactions.
• Non-renewable batteries stop working when reactants are used up.
• Rechargeable batteries can be recharged by reversing the chemical reaction with an external current.

"They contain chemicals that can produce a potential difference or voltage to power electrical appliances."

• Batteries generate voltage through chemical reactions to power devices.

"Rechargeable batteries can be recharged when a supplied current causes the reverse reaction to occur."

• Rechargeable batteries can be recharged by reversing the chemical reaction.

• Hydrogen fuel cells split water into hydrogen and oxygen through electrolysis.
• Recombining hydrogen and oxygen produces a voltage.
• Hydrogen fuel cells provide a renewable energy source.

"Hydrogen fuel cells work in a similar way water is split up into hydrogen and oxygen by electrolysis when they recombine a voltage is produced."

• Hydrogen fuel cells generate voltage by recombining hydrogen and oxygen, offering a renewable energy solution.