Atomic Structure, Isotopes and Mass Number

Structure of an Atom

• Atoms are the building blocks of everything, including DNA, cells, food, fragrances, and fuel.
• Atoms consist of subatomic particles: protons, neutrons, and electrons
• Protons are positively charged, located in the nucleus, and have a relative mass of 1 atomic mass unit (amu).
• Neutrons are neutral, located in the nucleus, and have a relative mass of 1 amu.
• Electrons are negatively charged, located in shells, and have a relative mass of 0.0005 amu.

Atomic Number

• Atomic number represents the number of protons in the nucleus of an atom.
• Each element is unique and based on its number of protons.

Mass Number

• Mass number represents the total number of protons and neutrons in the nucleus of an atom.
• Mass number = number of protons + number of neutrons.

Isotopes

• Isotopes are atoms of the same element with the same number of protons but a different number of neutrons.
• Isotopes have the same atomic number but different mass numbers.

Atomic Notation (Ion)

• Atomic notation includes the mass number, atomic number, chemical symbol, and ion charge to indicate gained or lost electrons.

Bohr Model

• The Bohr model consists of electrons located in specific levels or 'shells' orbiting the nucleus.
• Each shell has a maximum number of electrons, following the rule 2n^2 (where n = the shell number).
• Shells closer to the nucleus are occupied first due to lower energy.
• Shells fill in a specific order.
• The outer shell can never contain more than eight electrons.

Schrodinger Model

• Each 'shell' is composed of subshells.
• Each 'subshell' is composed of orbitals.
• Each 'orbital' can hold a maximum of two electrons.
• Electrons occupy the lowest possible energy levels (or subshells) first.

Aufbau Principle

• The Aufbau principle states that electrons fill the lowest energy shells or energy levels first.

Periodic Table

• Elements in the periodic table are arranged into vertical columns called groups.
• Elements in the same group have the same valence electrons.
• Horizontal rows are called periods; the period indicates the number of occupied electron shells.

Trends in the Periodic Table

• Core Charge is the electrostatic force of attraction of outer-shell (valence) electrons to the nucleus.
• Core charge = total number of protons – total number of inner-shell electrons.
• Atomic Radius measures the size of an atom from the center of its nucleus to its valence shell. Affected by core charge and number of occupied electron shells.
• Electronegativity is the ability of an atom to attract electrons to itself.
• The greater the electronegativity, the greater the ability of the atom to attract electrons towards its nucleus.
• First Ionisation Energy is the amount of energy required to remove the outermost electron from a gaseous atom. Moving across a period from left to right, there is a change from metals to metalloids and then to non-metals.
• Metallic character decreases across a period.
• Moving down a group, atomic radius increases, while electronegativity and first ionization energy decrease.
• Metals more easily lose electrons and become more reactive down a group.

Bonding

• Covalent Bonding involves the sharing of electrons between two non-metals to fill their valence shells.
• Sharing follows the octet rule, where atoms are most stable with eight electrons in their outer shell.
• More shared electron pairs result in a stronger covalent bond.
• Lewis structures (electron dot formulas or diagrams) do not include inner shell electrons.
• Valence shell electron pair repulsion theory (VSEPR) predicts the shapes of molecules based on electron pairs in the valence shell.
• Like charges repel, so negatively charged electron pairs are most stable when as far apart as possible.
• Lone pairs repel bond pairs more strongly than bond pairs repel each other.

Polarity

• Polar Covalent bonds form between two covalently bonded atoms with differing electronegativities.
• Non-Polar Covalent bonds form between two atoms with similar electronegativity.

Intramolecular Bonds

• Intramolecular Bonds, such as covalent bonds, occur within a molecule.
• Two factors affect the strength of the covalent bond: bond length and number of electrons shared.

Intermolecular Forces

• Intermolecular Forces are forces of attraction between molecules.
• Temporary Dipole-Dipole (dispersion forces) When electrons concentrate in one region, creating slight charges. These are very weak attractions.
• Permanent Dipole-Dipole occur between polar molecules when the partially positive end of one molecule attracts the partially negative end of another molecule.
• Hydrogen Bonds are a type of permanent dipole-dipole bond that occurs when hydrogen is bonded to nitrogen, oxygen, or fluorine. Highest melting and boiling points.
• Metallic Bonding occurs between two metals with low ionization energies.

Properties of Metallic Bonding

• Metals are arranged in a crystal lattice structure with valence electrons moving freely around metal cations.
• Metals can be bent and hammered into different shapes without breaking (malleability).
• Metals can be drawn into thin wires (ductility).
• Metals are good heat conductors due to delocalised electrons and metal ions; kinetic energy increases when heated.
• Metals are good electrical conductors due to freely moving, delocalised electrons.
• Metals have high melting and boiling points because of strong electrostatic forces.
• Metals appear lustrous because delocalized electrons readily reflect light.
• Metal reactivity describes how easily a metal loses electrons during a reaction.

Reactions of Metals

• Metals react with acids to produce a salt and hydrogen gas.
• Metals react with water to produce a metal hydroxide and hydrogen gas (typically Group 1 and 2 metals).
• Metals react with oxygen gas to produce a metal oxide.

This lesson explores the structure of atoms, including protons, neutrons, and electrons. It defines atomic number as the number of protons and mass number as the sum of protons and neutrons. It also explains isotopes as atoms of the same element with different numbers of neutrons.