Show Menu

GCSE AQA chemistry higher atomic structure Cheat Sheet by

made using the separate science GCSE specification.I couldn't find anything like this and thought everyone should have good resources for their GCSE's

atom,s­ymb­ols­,el­ect­ronic charge and isotopes

All substances are made of atoms.An atom is the smallest part of an element that can exist.
Compounds- are formed from elements by chemical reacti­ons.Co­mpounds contain two or more elements chemically combined in fixed propor­tions
Word equations- display chemical reactions, denoting reactants and products with their full chemical names.Example- Sodium hydroxide + hydroc­hloric acid ⟶ sodium chloride + water
Symbol equations- utilize the formulae of reactants and products to illustrate chemical reactions.Example- S + O₂ → SO₂
Half equations- specif­ically illustrate electron behavior during reactions where entities gain or lose electrons.Examples of Half Equations- Pb²⁺ + 2e⁻ → Pb 2Br⁻ → Br₂ + 2e⁻
Ionic equations- elucidate ion behavior during reacti­ons.Ex­ample of Ionic Equation Initial equation: HCl + NaOH → NaCl + H₂O Ionic equation: H⁺ + OH⁻ → H₂O
A mixture consists of two or more elements or compounds not chemically combined together.
Filtra­tion- Separates undiss­olved solid from a liquid­/so­lution mixture (e.g., sand from water). Utilizes a filter funnel equipped with filter paper placed over a beaker. The filter paper permits only liquid particles to pass, retaining solid particles as residue.
Crysta­lli­sation- Employed for separating a dissolved solid from a solution (e.g., copper sulphate from its aqueous solution). The solution is heated to create a saturated solution, followed by slow cooling to facilitate crystal growth. Crystals are harvested by filtering, washing with cold distilled water, and drying.
Simple Distil­lation- Separates a liquid and a soluble solid from a solution or pure liquid from a liquid mixture. The heating process initiates evapor­ation, producing vapor that condenses into pure liquid in a condenser. The remaining solid solute is left behind post complete liquid evapor­ation.
Fractional Distil­lation- Separates miscible liquids based on their boiling points (e.g., ethanol and water). The solution is heated to the boiling point of the lower boiling substance, which is then evaporated and collected separa­tely. Example: In a water-­ethanol mixture, ethanol (boiling point 78 °C) is evaporated first, followed by water (boiling point 100 °C).
Paper Chroma­tog­raphy- Separates substances with varying solubi­lities in a solvent (e.g., different dyes in black ink). A pencil­-drawn line on chroma­tog­raphy paper holds sample spots for analysis. The solvent ascends the paper via capillary action, carrying colored substances at different rates based on their solubi­lity. Results in separation of components and indicates purity or mixture status of a substance based on the number of spots developed.
The discovery of the electron led to the plum pudding model of the atom. The plum pudding model suggested that the atom is a ball of positive charge with negative electrons embedded in it.
The results from the alpha particle scattering experiment led to the conclusion that the mass of an atom was concen­trated at the centre (nucleus) and that the nucleus was charged. This nuclear model replaced the plum pudding model.
proton-+1 charge electron--1 neutron- 0
The number of protons in an atom of an element is its atomic number. All atoms of a particular element have the same number of protons. Atoms of different elements have different numbers of protons.
mass number, symbolized as A, is the sum of protons and neutrons in an atom's nucleus.
Isotopes- are atoms of the same element but with different numbers of neutrons.Number of neutrons (n) = mass number - atomic number
Relative atomic mass equation-∑ isotope mass x isotope abundance / 100
The electrons in an atom occupy the lowest available energy levels (innermost available shells). The electronic structure of an atom can be repres­ented by numbers or by a diagram.

The Periodic Table

The elements in the periodic table are arranged in order of atomic (proton) number and so that elements with similar properties are in columns, known as groups. The table is called a periodic table because similar properties occur at regular intervals.
Elements in the same group in the periodic table have the same number of electrons in their outer shell (outer electrons) and this gives them similar chemical proper­ties.
Initial Ordering Methods- Before subatomic particles were discov­ered, elements were arranged by atomic weight, not atomic number. Patterns began to emerge when elements were organized by mass, leading to the term 'perio­dic'. Some elements were forced into positions to fill gaps, while others were incorr­ectly placed based on atomic weight only.
Mendel­eev's Contri­bution- First Draft in 1869 Russian chemist Dmitri Mendeleev made the first draft of the periodic table in 1869. Elements were organized in vertical columns based on their properties and compound charac­ter­istics. Horizontal Arrang­ements and Patterns As Mendeleev arranged elements by increasing atomic weight, chemically similar elements naturally fell into the same columns. There were some exceptions where elements didn't follow this pattern. Innovation and Predic­tions Mendeleev did not force elements into specific positions, he left gaps for undisc­overed elements. He even switched the positions of elements to maintain property consis­tency. Mendeleev used existing element properties to predict the charac­ter­istics of undisc­overed elements, like "­eka­-si­lic­on" now known as germanium.
Mendel­eev's Limita­tions- Mendeleev had no knowledge of isotopes, leading to some inaccu­racies. He did consider both atomic mass and chemical properties when sorting, but inaccu­racies remained. Impact of Subatomic Particles Once subatomic particles were discov­ered, atomic numbers were calculated for each element. The modern Periodic Table uses atomic numbers, aligning with Mendel­eev's original patterns.
Elements that react to form positive ions are metals. Elements that do not form positive ions are non-metals
The elements in Group 0of the periodic table are called the noble gases. They are unreactive and do not easily form molecules because their atoms have stable arrang­ements of electrons. The noble gases have eight electrons in their outer shell, except for helium, which has only two electrons. The boiling points of the noble gases increase with increasing relative atomic mass (going down the group)
The elements in Group 1 of the periodic table are known as the alkali metals and have charac­ter­istic properties because of the single electron in their outer shell. Students should be able to describe the reactions of the first three alkali metals with oxygen, chlorine and water. In Group 1, the reactivity of the elements increases going down the group.
The elements in Group 7 of the periodic table are known as the halogens and have similar reactions because they all have seven electrons in their outer shell. The halogens are non-metals and consist of molecules made of pairs of atoms. Students should be able to describe the nature of the compounds formed when chlorine, bromine and iodine react with metals and non-me­tals. In Group 7, the further down the group an element is the higher its relative molecular mass, melting point and boiling point. In Group 7, the reactivity of the elements decreases going down the group. A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt

Reactions of Group 1+Trends and Properties

Reactions with Water- Alkali metals react more vigorously with water as you move down the group. They are generally stored in oil to prevent reactions with air and water vapor.
Reactions with Oxygen- Alkali metals react with oxygen to form metal oxides that cause tarnis­hing.
Reactions with Chlorine- All alkali metals react intensely when heated with chlorine gas, forming metal chlorides. The reaction gets more vigorous as you move down the group.
Softness and Density- The metals get softer as you move down the group, with potassium being the exception which has lower density than sodium. The first three metals in this group are less dense than water.
Melting Points- Melting points for these metals decrease as you move down the group.
Reacti­vity- Reactivity increases down the group, with atoms needing to lose just one outer electron to attain noble gas config­ura­tion. As the number of shells increases down the group, the outer electron is farther from the nucleus and thus more easily lost, increasing reacti­vity.

Comparing Transition Metals and Group 1 Elements

Position in Periodic Table- Transition elements are found between Groups 2 and 3 in the center of the periodic table. They exhibit the typical metallic properties but have key differ­ences compared to Group 1 metals.
Charge on Ions- All Group 1 metals form ions with a +1 charge. Transition metals can form ions with variable charges, like Fe2+ and Fe3+ ions in the case of iron.
Physical Proper­ties- Transition metals are much harder, stronger, and denser compared to the soft and light Group 1 metals. They have signif­icantly higher melting points. For example, titanium melts at 1,688 ºC, while potassium melts at 63.5 ºC.
Reacti­vity- Transition metals are less reactive than Group 1 metals. Alkali metals (Group 1) react rapidly with water, oxygen, and halogens. Transition metals react more slowly or may not react at all.
Reactivity with Oxygen- Group 1 metals tarnish quickly in the presence of oxygen, forming metal oxides. Iron, as a transition metal, takes several weeks to react with oxygen to form iron oxide (rust), and it needs water for this reaction.

Transition Metals: Properties and Applic­ations

Properties of Transition Metals- Most known metals are transition metals, exhibiting typical metallic proper­ties. These metals are lustrous, hard, strong, and good conductors of heat and electr­icity. Transition metals are dense and have high melting points. They can have multiple oxidation states, allowing them to lose different numbers of electrons based on their chemical enviro­nment. Variab­ility in Compounds Compounds with transition elements in different oxidation states have varying properties and colors when dissolved in water.
Applic­ations of Transition Metals
Catalysis Transition metals are widely used as catalysts, substances that speed up chemical reactions without being consumed. Their catalytic properties are due to their ability to switch between multiple oxidation states. They form complexes with reagents, facili­tating electron donation and acceptance within a chemical reaction.
Common Catalysts Iron is utilized in the Haber Process. Vanadium pentoxide (V2O5) is used in the Contact Process for sulfuric acid produc­tion. Nickel is used for hydrog­enating alkenes.
Medicine- Transition metals find applic­ations in medicine, especially in limb and joint replac­ements. Titanium is signif­icant here, as it can bond with bones due to its high biocom­pat­ibi­lity.
Other Industrial Applic­ations- These metals are used in the creation of colored compounds for dyes, paints, and other applic­ations. Additional uses include making stained glass, crafting jewelry, and in anti-c­orr­osive materials.


No comments yet. Add yours below!

Add a Comment

Your Comment

Please enter your name.

    Please enter your email address

      Please enter your Comment.

          Related Cheat Sheets