Chemistry OCR B EL key definitions

Ammonium Ion

An inorganic ion with the formula NH₄⁺.

Amount of Substance

The quantity of a chemical species, measured in moles. Used as a way of counting atoms. The amount of substance can be calculated using:

Number of moles = Mass/Mr

Number of moles = (Pressure*Volume)/(Gas constant*Temperature)

Number of moles = Concentration*Volume

Atomic Number

The number of protons in the nucleus of an atom

Atomic Orbital

A region of space around the nucleus that can hold up to 2 electrons with opposite spins. There is one orbital in the s subshell, 3 orbitals in the p subshell and 5 orbitals in the d subshell. Orbitals are filled in order of increasing energy, with orbitals of the same energy being occupied singly before pairing.

Avogadro Constant (NA)

The number of particles per mole of substance (6.02×10²³ mol^-1)

Carbonate

An ion with the formula CO₃^2-.

Composition by Mass

The relative mass of each element in a compound.

Concentration

The amount of substance per unit volume. Units are given in g/dm³ or mol/dm³.

Electron

A negatively charged subatomic particle that orbits the nucleus at varying energy levels. The relative mass of an electron is 1/1840.

Electronic Configuration

The arrangement of electrons into orbitals and energy levels around the nucleus of an atom/ion. E.g. Ca: 1s²2s²2p⁶3s²3p⁶4s²

Empirical Formula

The simplest whole-number ratio of atoms of each element present in a compound.

Energy Level

The shell that an electron is in.

Fusion Reactions

Lighter nuclei join to give heavier nuclei (under conditions of high temperature and pressure). This is how certain elements are formed.

Hydroxide

An ion with the formula OH^-.

Ion

Particle formed when an atom/molecule loses or gains electrons. This gives it an overall charge - a positive charge if it has lost at least one electron and a negative charge if it has gained at least one electron.

Ionic Compound

A compound made up of oppositely charged ions that are held together by electrostatic forces.

Isotopes

Atoms of the same element with the same number of protons and electrons but different numbers of neutrons. Isotopes of an element have different mass numbers.

Mass Number

The total number of protons and neutrons in an atom's nucleus.

Molecular Formula

Formula which shows the actual number and type of atoms of each element in a molecule.

Neutron

A neutral subatomic particle found in the nucleus of an atom. The relative mass of a neutron is 1.

Nitrate

An ion with the formula NO₃^-.

p Orbital

A dumbbell-shaped region in which up to two electrons can be found. There are three p orbitals at right angles to each other, so in total, the p subshell can hold up to 6 electrons.

Percentage Yield

The percentage ratio of the actual yield of the product from the reaction compared with the theoretical yield.

Percentage yield = (Actual yield/ Theoretical yield)*100

Proton

A positively charged subatomic particle found in the nucleus of an atom, contributing to the atomic number. The relative mass of a proton is 1.

Relative Atomic Mass

The weighted mass of an atom compared with 1/12th mass of an atom of carbon-12.

Relative Formula Mass

The mass of the formula unit of a compound with a giant structure. For example, NaCl has a relative formula mass of 58.44g mol^-1.

Relative Isotopic Mass

The mass of an atom of an isotope compared with 1/12th mass of an atom of carbon-12.

Relative Molecular Mass

The average mass of one molecule of an element or compound compared to 1/12th the mass of an atom of carbon-12.

s Orbital

Spherical and symmetrical regions around the nucleus, they can each hold up to two electrons. They are the lowest energy orbitals in an atom.

Shell

The energy level that an orbital is in around the nucleus of an atom. The shell closest to the nucleus is the first shell. The outermost shell that is occupied by electrons is the valence shell. Shells are organized by increasing energy levels, with each shell capable of holding a specific number of electrons.

Standard Solution

A solution that has a known concentration of a compound/element. It is used in titrations to determine the concentration of unknown solutions.

State Symbols

Symbols which show the physical state of the substance during the reaction, they are usually in brackets: gas (g), liquid (l), solid (s) and aqueous (aq). Aqueous means the substance is dissolved in water.

Sub-Shell

A subdivision of the electronic shells into different orbitals. The types of subshell are s, p, d and f.

Sulfate

An ion with the formula SO₄^2-.

Titration

An experimental technique used to determine the concentration of an unknown solution by using a second solution with a known concentration. This is done by adding the known solution to the unknown solution until the reaction reaches a specific endpoint, often indicated by a color change.

Water of Crystallisation

Water molecules that form part of the crystalline structure of a compound. These water molecules are essential for maintaining the stability and integrity of the crystal lattice.

Covalent Bond

The strong electrostatic attraction between two nuclei and the shared pair of electrons between them. Polar covalent bonds occur when there is asymmetric electron distribution within the covalent bond due to a difference in electronegativities.

Covalent Substance

A substance that is made up atoms that are covalently bonded to each other. These substances typically have low melting and boiling points and may exist as gases, liquids, or solids at room temperature.

Dative Covalent Bonding

Occurs when one atom donates both electrons in a covalent bond. This type of bonding typically involves a donor atom and an acceptor atom, resulting in a bond where both electrons come from the same atom. For example, in NH₄ for one of the N-H bonds, nitrogen provides both of the bonding electrons.

Dot-and-Cross Diagram

Diagrams used to model the bonding that occurs in a simple molecule. The shells of an atom are drawn as circles, with crosses or dots marked on the circles to represent the electrons. The circles overlap when there is a covalent bond. The electrons from one atom are drawn as dots, and the electrons from a different atom as crosses. These diagrams visually represent how electrons are shared between atoms in covalent bonding.

Electrical Conductivity

A measure of the amount of electrical current a material can carry or its ability to carry the current. It is influenced by the presence of free-moving charged particles, such as ions or electrons, in the material.

Electron Pair Repulsion

The repulsion between pairs of electrons which mean the shape that a molecule adopts has the pairs of electrons positioned as far apart as possible. As a result, carbon atoms in alkanes have a tetrahedral shape and a bond angle of 109.5 degrees.

Electrostatic Attraction

The attraction between two species with opposite charges. This force is responsible for holding ionic compounds together and influencing molecular structures.

Giant Atomic Structure

Large structures containing lots of atoms that are covalently bonded to each other, they are usually arranged in a regular lattice. E.g. Diamond. These structures have high melting and boiling points due to the strong covalent bonds that must be broken to change their state.

Giant Ionic Lattice

A regular repeating structure made up of oppositely charged ions. These lattices are typically found in ionic compounds and exhibit high melting and boiling points due to the strong electrostatic forces between the ions.

Ionic Bond

Strong electrostatic attraction between two oppositely charged ions. The strength of attraction depends on the relative sizes and charges of ions. Ionic bonds are formed when one atom donates an electron to another, resulting in the formation of positively and negatively charged ions that attract each other.

Ionic Charge

The electrical charge of an ion caused by the gain (negative charge) or loss (positive charge) of electrons. The magnitude of the charge is related to how many electrons have been lost or gained as electrons have a relative charge of -1.

Ionic Compound

A compound made up of anions and cations held together by ionic bonds, which arise due to the electrostatic attraction between oppositely charged ions. These structures are neutral overall.

Linear

The shape of a molecule when the central atom has 2 bonding pairs and no lone pairs of electrons. In this arrangement, the atoms are positioned at 180 degrees to each other.

Melting Point

The temperature at which the element changes from a solid state to a liquid state.

Metallic Bonding

Strong electrostatic attraction between positive metal ions and the sea of delocalised electrons that surround them. This bonding allows metals to conduct electricity and heat, and gives them malleability and ductility.

Non-Linear/bent

The shape of a molecule when the central atom has 2 bonding pairs and 2 lone pairs.

Octahedral

The shape of a molecule when the central atom has 6 bonding pairs. This geometry results in a symmetrical arrangement of bonds, typically with 90-degree angles between them.

Pyramidal

The shape of a molecule when the central atom has 3 bonding pairs and 1 lone pair. This results in a trigonal pyramidal shape, where the bond angles are approximately 107 degrees.

Simple Molecular Structure

Atoms that are covalently bonded together to form relatively small molecules. These molecules have low melting and boiling points and are often gases or liquids at room temperature.

Solubility In Water

The degree to which a substance can dissolve in water at a certain temperature. It is influenced by factors such as temperature, pressure, and the nature of the solute and solvent.

Tetrahedral

The shape of a molecule when the central atom has 4 bonding pairs. It has bond angles of 109.5 degrees.

Trigonal Bipyramidal

The shape of a molecule when the central atom has 5 bonding pairs. It features two distinct types of bond angles: 90 degrees between the equatorial and axial positions, and 120 degrees between the equatorial positions.

Trigonal Planar

The shape of a molecule when the central atom has 3 bonding pairs. It has bond angles of 120 degrees.

Anion

A negatively charged ion, formed when an atom gains at least one electron, e.g. S^2-. Anions are typically more stable in ionic compounds and play a crucial role in various chemical reactions.

Cation

A positively charged ion, formed when an atom loses at least one electron, e.g. Na⁺. Cations are essential in ionic bonding and often participate in chemical reactions.

Charge Density

The ratio of the charge of an ion compared to its volume, for example, a 3+ ion will have a higher charge density than a 1+ ion of similar size. It influences the strength of ionic bonds and the properties of ionic compounds.

First Ionisation Energy

The energy required to remove 1 mole of electrons from 1 mole of gaseous atom to form one mole of gaseous 1⁺ ions. For example, Mg_(g) → Mg⁺_(g) + e^-. It reflects the atom's ability to hold onto its electrons and varies across the periodic table.

Precipitation Reaction

A reaction in which solutions react to form an insoluble product. When combined with acidified silver nitrate, halide ions react to form different coloured precipitates depending on the ion present. The colour of the precipitate formed can be used to identify which halide is present. This process is commonly used in qualitative analysis to detect specific ions in solution.

Precipitate

The solid formed from a reaction in solution. It occurs when the concentration of a dissolved substance exceeds its solubility, resulting in the formation of an insoluble compound.

p-block Element

Elements found in groups3-8/0 of the periodic table. p-block non-metals generally undergo reduction reactions. They include elements such as carbon, nitrogen, oxygen, and the noble gases, which have their highest energy electrons in the p orbital.

s-block Element

Elements in group 1 and 2 of the periodic table. s-block elements generally undergo oxidation reactions. They include alkali and alkaline earth metals, which have distinctive properties such as high reactivity and the ability to form cations.

Alkali

A base which is soluble in water. Alkalis produce hydroxide ions (OH-) in solution and can neutralize acids.

Brønsted-Lowry Acid

Proton donors. These species release hydrogen ions in solution. They can be either strong or weak acids depending on their ability to donate protons.

Brønsted-Lowry Base

Proton acceptors. These species can accept hydrogen ions in solution, often leading to the formation of conjugate acids.

Diprotic Acid

An acid that can release two H⁺ ions upon dissociation, e.g. H₂SO₄. They can donate protons in two steps, resulting in a diprotic equilibrium.

Monoprotic Acid

An acid that can release only one H⁺ ion upon dissociation, such as in HCl.

Neutralisation

A reaction between an acid and a base which react together to form water and a salt. The ionic equation for neutralisation is:

H⁺_(aq) + OH^-_(aq) → H2O_(l)

Polyprotic Acid

An acid that can release more than two H⁺ ions upon dissociation , e.g. H³PO⁴. They can donate protons in multiple steps, resulting in polyprotic equilibria.

Strong Acid

An acid that dissociates/ionises almost completely in water. This means nearly all the H⁺ ions will be released. Examples include HCl and H²SO⁴. Strong acids have a low pH and are highly reactive.

Strong Base

A base that dissociates/ionises almost completely in water. This means nearly all the OH^- ions will be released, resulting in a high pH and increased reactivity. Examples include NaOH and KOH.

Weak Acid

An acid that only partially dissociates in water, meaning only some of the H⁺ ions are released. This results in a higher pH compared to strong acids and includes examples like acetic acid (CH³COOH) and citric acid. Weak acids are characterized by their ability to establish an equilibrium between the undissociated acid and the ions in solution.

Weak Base

A base that partially dissociates in water, releasing fewer OH^- ions compared to strong bases. This results in a lower pH and includes examples such as ammonia (NH³) and bicarbonate (HCO^3-). Weak bases also establish an equilibrium between the undissociated base and the ions in solution.

Absorption Spectra

A spectrum of frequencies of electromagnetic radiation that has been transmitted through an atom or molecule, that shows dark bands due to the absorption of the radiation at those specific wavelength.

Electromagnetic Spectrum

The range of frequencies of electromagnetic radiation and the respective wavelengths. It includes gamma rays, X-rays, ultraviolet, visible light, infrared, microwaves, and radio waves.

Emission Spectra

A spectrum of frequencies of electromagnetic radiation that has been emitted by an atom or molecule undergoing a transition from a state with higher energy to a state with lower energy. It shows bright lines or bands at specific wavelengths corresponding to the energies of the emitted photons.

Flame Test

An analytical technique used to identify certain elements and ions based on the colour produced when a nichrome wire is dipped into a solution of the species and held in a blue Bunsen flame to produce characteristic colours.

Infrared

The part of the electromagnetic spectrum that has wavelengths between 780nm and 1mm. It is primarily associated with heat and is used in various applications such as thermal imaging and spectroscopy.

Ultra-Violet

The portion of the electromagnetic spectrum with wavelengths shorter than visible light, typically ranging from 10nm to 400nm. It is known for its ability to cause fluorescence and is used in applications such as sterilization and phototherapy.

Visible Light

The part of the electromagnetic spectrum that is visible to the human eye, with wavelengths ranging from approximately 380nm to 700nm. It encompasses all colours perceived as light, including red, orange, yellow, green, blue, indigo, and violet.

Mass Spectrometry

A technique used to identify compounds and determine their relative molecular mass and the relative abundance of isotopes in a sample. It involves ionizing chemical species and sorting the ions based on their mass-to-charge ratio.

Relative Abundance

The amount of substance compared with another substance in a mixture, often expressed as a percentage or ratio.

Relative Abundance (of Isotopes)

The percentage of atoms found within a naturally occurring sample of an element that has a specific atomic mass. It is used to determine the distribution of isotopes in a sample and is crucial for calculating the average atomic mass of an element.

Relative Atomic Mass

The weighted mean mass of an atom compared with 1/12th mass of an atom of carbon-12. It reflects the average mass of an element's isotopes, accounting for their relative abundances.