Structure of Matter

Flashcards reviewing the properties and structure of matter, atomic structure, periodicity, and bonding, based on lecture notes.

Physical Properties

The characteristics of a substance that can be observed or measured without changing it into a different substance.

Chemical Properties

The properties of a substance associated with the chemical changes it undergoes when mixed with other substances, heated, or exposed to light, resulting in a change of the substance into a different substance.

Homogeneous Substances

Substances with uniform composition throughout, such as pure water, air, or petrol.

Heterogeneous Substances

Substances with non-uniform composition, such as concrete.

Melting Point

The lowest temperature at which a solid changes to a liquid.

Boiling Point

The lowest temperature at which a liquid boils (changes from liquid to a gas).

Density

Defined as mass per unit volume (d=m/v).

Gravimetric Analysis

Analysis by mass to determine the quantities (mass) of substances present in a sample.

Element

Made of just one type of atom.

Monoatomic elements

Made up of only one atom.

Compounds

Contain different types of atom chemically bonded together in definite proportions.

Periodic Law

The properties of the elements vary periodically with their atomic numbers.

Metals

Shiny and reflect light; malleable; ductile; silver-colored; dense; high melting point; high boiling point; high tensile strength; good conductors of electricity; good conductors of heat.

Non-metals

Not malleable; not ductile; dull in colour, not shiny; not dense; lower in melting and boiling points; poor conductors of electricity; poor conductors of heat.

Metalloids

Have some metallic and some non-metallic properties.

Protons

Positively charged particles with a mass of approximately 1.673 × 10-27 kg.

Neutrons

Particles with no charge and a mass of approximately 1.675 × 10-27 kg.

Isotopes

Atoms of the one element that have different numbers of neutrons in the nucleus (same atomic number but different atomic mass number).

Radioisotopes

Isotopes that spontaneously emit radiation.

Radioactivity

Spontaneous emission of radiation.

Alpha (𝝰) particles

A helium nucleus, charge +2, mass 4, low penetrating power.

Beta (𝞫) particles

Electron, charge -1, mass 1/2000, moderate penetrating power.

Gamma (ɣ) radiation

Electromagnetic radiation, charge 0, mass 0, high penetrating power.

Rules to determine electron configuration

Each shell can contain a maximum number of electrons; lower energy shells fill before higher energy shells; electron shells fill in a particular order.

Orbital

The volume of space surrounding the nucleus of an atom through which one or two electrons may randomly move.

Subshells

Energy levels within an atom.

Pauli Exclusion Principle

Each orbital can contain a maximum of two electrons, with each electron having a different spin.

Aufbau principle

The lowest energy orbitals are always filled with electrons first.

Hund's rule

Every orbital in a subshell must first be filled with one electron with the same spin before an orbital is filled with a second electron.

Relative Isotopic Masses

The individual isotopes of each element have a relative isotopic mass (Ir). The relative isotopic mass of an isotope is the mass of an atom of that isotope relative to the mass of an atom of Carbon-12, taken as 12 units exactly.

Relative Isotopic Abundance

The percentage of that isotope in the naturally occurring element.

Relative Atomic mass (Ar)

The weighted average of the relative masses of the isotopes of the element on the 12C scale.

Relative Molecular mass (Mr)

Equual to the sum of the relative atomic masses of the atoms (from the periodic table) in the molecule.

Relative Formula mass

The mass of a formula unit and is calculated by taking the sum of the relative atomic masses of the elements in the formula.

Atomic Emission Spectroscopy

Heating an element can cause an electron to absorb energy and jump to a higher energy level. Subsequently, it returns to the lower energy level, releasing light.

Flame Test

Metallic elements produce a characteristic color when placed in a Bunsen flame.

Blocks of the Periodic Table

The subshell of an element's block on the periodic table where the element's highest energy electrons are located.

Core Charge

A measure of the attractive force felt by the valence shell electrons towards the nucleus.

Atomic Radius

The distance from the nucleus to the valence shell electrons.

First Ionisation energy (IA)

The energy required to remove one electron from an atom of an element in the gas phase. M(g) → M+ + e-

Electronegativity

The ability of an atom to attract electrons towards itself.

Ionic compounds

Made by the chemical combination of metallic and non-metallic elements.

Ionic bonding

The chemical bonding that involves the outright transfer of electrons from one atom to another; bonding that consists of electrostatic attraction between the positive and negative ions formed by this transfer of electrons.

Single Covalent Bond

The covalent bond formed when atoms share two electrons, one from each atom.

Lewis Electron Dot Diagram

Show the valence shell electrons of an atom because only these electrons are involved in bonding.

Covalent Molecular Substances

Molecules have strong covalent bonds inside, and the much weaker bonds between molecules.

Intramolecular bonds

Bonds within a molecule.

Intermolecular bonds

Bonds between molecules.

Allotropes

Some elements can exist with their atoms in several different structural arrangements.

Valence Shell Electron-Pair Repulsion (VSEPR) Theory

Uses knowledge of the valence electrons in the atoms of a molecule to predict the shape of the molecule. The VSEPR theory is based on the principle that negatively charged electron pairs in the outer shell of an atom repel each other.

Covalent Network structures

Solids in which the covalent bonding extends indefinitely throughout the whole crystal. Formulas are written as the simplest ratio of the elements present in the network.

Fullerenes

Molecules containing a spherically arranged group of carbon atoms in a series of pentagons and hexagons, similar to the shape of a soccer ball.

Nanotubes

Has a long, hollow structure with walls formed from graphene.

Metallic Bonding Model

Where positive ions or cations are arranged in a closely packed three-dimensional network structure, or lattice, and negatively charged electrons move freely throughout the lattice.

Electrical Conductivity, Aluminum

The sea of delocalized electrons can move through the lattice towards a positively charged electrode.

Non-polar diatomic molecules

When two atoms form a covalent bond, you can regard the atoms as competing for the electrons being shared between them. If the two atoms in a covalent bond are the same (i.e. have identical electronegativities), then the electrons are shared equally between the two atoms.

Bond Types

A great enough electronegativity imbalance that will transfer electrons between two atoms, resulting in electron transfer and the formation of ionic bonds.

Dispersion forces

The forces of attraction between non-polar molecules caused by temporary dipoles in the molecules that are the result of random movement of the electrons surrounding the molecule.

Intermolecular forces

dipole–dipole forces; hydrogen bonding; dispersion forces