Hitting the jackpot

Dalton's atomic theory

1. all matter is made of tiny particles called atoms.
2. the atoms of an element are always the same (mass, shape, size, etc) while the atoms of different elements are different.
3. compounds form when atoms combine (different elements), atoms combine in whole/consistent (always a fixed specific #) number ratios.
4. reactions involve the combination of atoms, the atoms themselves do not change and are not destroyed/created.

JJ thomson and his experiment

• Observed that the rays were deflected by electric and magnetic fields

Conclusion:

• Rays are made of negatively charged particles → electrons.

• Led to the idea that atoms are not indivisible.

• Proposed the plum pudding model: electrons embedded in a positive “soup.”

Plum pudding model

The atom is a sphere of diffuse positive charge with electrons embedded throughout, like plums in a pudding (or pie).

Ernest rutherford and the gold foil experiment

• Fired alpha particles (2+) at thin gold foil; detected with a fluorescent screen.

• Most passed straight through → atom is mostly empty space.

• Some deflected → presence of a concentrated positive charge.

• Very few (~1 in 8000) bounced back → tiny, dense, positively charged nucleus.

Niels Bohr

electrons orbit the nucleus somewhat like planets orbit the sun --> planetary model. Has energy levels

atomic number

number of protons in the nucleus defining the element

Relative Atomic Mass

Relative atomic mass is the weighted average mass of the atoms of an element compared with 1/12 of the mass of one atom of carbon-12, taking into account the natural abundances of its isotopes.

Isotopes

Same number of protons but different number of neutrons

Properties of isotopes

Only physical properties change (boiling, melting, density) not chemical.

atomic mass.

Actual mass of a single atoms measured in amu, based on nucleon number

Mass spectra

It separates atoms/ions based on the mass and charge (m/z) to tell u relative isotopic mass, abundance of each isotope, and the Ar of each element.

1. A sample put into the spectrometer and vaporised;

2. An electron beam creates ions with a 1+ charge

3. Ions are accelerated by an electric field and focused

4. Ions are deflected by an electromagnet. Changing the strength of the electromagnetic changes the deflecting pattern based on ion mass

5. The ions are detected by the detector and the mass spectrum plotted

continuous spectrum

has all wavelengths of light.

Hydrogen emission spectrum

When an excited electron falls to a lower energy level, hydrogen emits light at specific colours. There are four discrete visible lines that converge at higher levels (Balmer series) are: red (3→2), blue-green (4→2) , blue-violet (5→2) and violet (6→2).

orbitals

Orbitals in atoms are regions around the nucleus where there is a high probability of finding an electron. They come in different types with distinct shapes: s orbitals are spherical and one exists per energy level, p orbitals are dumbbell-shaped with three per energy level starting from n = 2, d orbitals are clover-shaped with five per level starting from n = 3, Each orbital can hold a maximum of two electrons with opposite spins.

Aufbau's principle

states that the order in which the atomic orbitals of atoms are filled with electrons is the order of increasing energy

Pauli's exclusion principle

no two electrons in an atom can have the same set of four quantum numbers: the Pauli principle states that a maximum of two electrons can occupy an atomic orbital – and these electrons must have opposite spins

Hund's rule

Electrons are added singly first to an orbital in the same energy level with parrallel spins. Then the second electron is added to each orbital

exceptional electron configurations

Cr and Copper. They have exceptional configurations as they take an electron form the four s shell to have a half filled or fully filled 3d shell. This is because having this half or fully filled 3d shell provides extra stability

trends on first IE

Going down a group first IE decreases due to electron shielding and the outermost electrons being farther away from the nucleus. Across a period first IE increases as there are more protons, thus more charge attracting the electrons

radioisotope

Unstable form of an atom (too many or too little neutrons) resulting in it being radioacive

Atoms divisibility

They are made up of smaller units protons and neutrons which can leave the atom in alpha decay. First IE, electrons can leave the nucleus to form an ion

First IE

the energy required to remove one mole of electrons from one molecule isolated gaseous atoms to form one mole of gaseous 1+ ions under standard

thermodynamic conditions

Line spectrum

an emission spectrum thathas only certain wavelengths or frequencies of visible light (it takes the form of a series of bright lines on a back background)

Ionization

An atom gaining or losing electrons becoming a charge element (cation or anion)

Ionic compounds

Made up of a cation and anion, electrons are not shared, overall charge must be neutral. Ionic compounds have high melting and boiling points because oppositely charged ions are held together by strong electrostatic attractions in a giant crystal lattice. They are hard due to these strong forces, but brittle because when layers shift under force, like charges repel and the lattice breaks.

Electrical conductivity of ionic compounds:

Ionic compounds do not conduct electricity when solid because ions are fixed in place. They conduct when molten or dissolved in water because ions are free to move and carry charge.

Hydration:

When an ionic compound dissolves in water, water molecules surround and stabilize the ions.

Lattice enthalpy:

The energy released when one mole of a solid ionic compound is decomposed to form gaseous ions. Exothermic when going from gas to solid. The more energy released the more stable the lattice

Factors affecting lattice enthalpy

Higher ionic charges lead to stronger electrostatic attraction between ions. Charge has a greater effect on lattice enthalpy than ionic size.

Smaller ions allow ions to be closer together, increasing attraction. Smaller ionic radius results in higher lattice enthalpy.

Metallic bonding

Metallic bonding consists of positive metal ions arranged in a lattice surrounded by a sea of delocalized electrons, is non directional. The attraction between the ions and electrons holds the metal together.

Factors affecting strength of metallic bonding

The most important factor is the number of delocalized electrons. More delocalized electrons increase the electrostatic attraction between the positive metal ions and the electron sea, strengthening metallic bonding. Cation size is secondary: smaller cations strengthen the attraction, but electron number has the greater effect overall.

Metallic bonding down a group:

As atomic radius increases down a group as number of energy shells and electrons increases. Weakening metallic bonding as the electron and positive metal ions cannot be as close to each other. This causes melting points to decrease.

Metallic bonding across a period:

Across a period, nuclear charge increases while atomic radius remains similar. This strengthens metallic bonding and increases melting and boiling points.

Malleability:

Malleability is the ability of a material to be hammered or compressed into shape without breaking.

Ductility:

Ductility is the ability of a material to be stretched into a wire without fracturing.

Atomic radius:

Atomic radius is defined as half the distance between the nuclei of two bonded atoms in a diatomic molecule.

Factors affecting atomic radius:

Atomic radius depends on energy level, nuclear charge, and shielding. Higher energy levels and electron shielding increase radius while higher nuclear charge decreases radius.,

Ionic radius

Ionic radius is the radius of the most common ion formed by an atom.

Cation size:

Cations are smaller than their parent atoms because losing electrons reduces electron–electron repulsion and allows electrons to move closer to the nucleus.

Anion size:

Anions are larger than their parent atoms because gaining electrons increases electron–electron repulsion, causing the electron cloud to expand.

Electronegativity:

Electronegativity is a measure of an atom’s ability to attract a bonding pair of electrons. It decreases down a group as atomic size increases and bonding electrons are further from the nucleus.¡

Electron affinity:

Electron affinity is the energy released when one mol of electron is added to a mol neutral gaseous atom to form one mol of gaseous 1-ion . It generally increases across a period and decreases down a group, but varies depending on electron configuration stability. The more energy released the more reactive an element is.

Alloys:

Alloys are mixtures of metals designed to improve properties such as strength, hardness, or resistance to corrosion. Are less malleable as the distorted structure compared to the pure metal makes it harder for the close-packed layers of cations to slide over each other. Thus also making them more brittle.

Trends of Melting point of halogens

As it goes down the group melting point increase because as atomic size and electron count increases London dispersion forces become stronger thus making it harder to separate the diatomic molecules

Trends of Alkali earth metals reactivity

Reactivity increases down a group because the distance from the nucleus to the valence electrons increases as you go down the group, the energy needed to remove that electron gets lower. This results in the reactivity increases as you go down a group.

Halogen reactivity down a group

Going down the group electron affinity decreases due to increased nuclear charge being offset by increased electron shielding and more energy levels thus become more non metals and less reactivity.

Oxidation

Loss of e^-, increased in oxidation state and loss of hydrogen. Oxidation does not occur without reduction. Species that is oxidized is called the oxidizing agent. The easier an atom to be oxidised the easier to lose electrons. As you go down group 1 the ionization energy decreases which indicates that electrons are lost more readily.

Reduction

When a species gains electrons, decrease in oxidation state or gain of hydrogen. The easier it is for an atom to gain electrons the more likely it is to be reduced. As you go down group 17 the electron affinity decreases which indicates that electrons are not gained as easily

Amphoteric

Can act both as a acid and a base.

Half equations

Involves only one element and shows the number of electrons gained or lost by that element

Transition elements

Transitional elements is an element with a partially filled d-orbital or has a common oxidation state with an incomplete d-orbital, then it is known as a transition element. However, zinc is not classified as a transition metal as it only has one stable ion with an oxidation state of 2+

Alloys

An alloy is a mixture of two or more elements, where at least one element is a metal.

Are less malleable as the distorted structure compared to the pure metal makes it harder for the close-packed layers of cations to slide over each other.

Acid

pH lower than 7 formed by non metals.

Base

pH higher than seven formed by alkali and alkali earth metals.

Ionic bonding:

Ionic bonding involves one element gaining electrons while the other loses electrons, in combination the compound must have a neutral charge.

Redox Reaction

Transfer of electrons between two species, thus one species is oxidized and one is reduced

Net ionic equation

A net ionic reaction involves only substance that are oxidized or reduces, thus the sum of the half equations

Oxidizing agent:

Leads to the oxidation of other elements, and itself is reduced

Reducing agent

Leads to the reduction of other elements and itself is oxidized

Lattice Enthalpy of formation

The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states, under standard conditions.

Sodium

Na₂O

Giant ionic lattice

Basic

Reacts with water to form NaOH(aq). Neutralises acids.

Magnesium

MgO

Giant ionic lattice

Basic (weak)

Reacts slowly with water to form Mg(OH)₂. Reacts with acids to form Mg²⁺ salts.

Aluminium

Al₂O₃

Giant ionic lattice (high charge density)

Amphoteric

No reaction with water. Reacts with acids → Al³⁺ salts. Reacts with bases → aluminate ions.

Silicon

SiO₂

Giant covalent network

Acidic

No reaction with water. Reacts with strong bases → silicates. No reaction with acids (except HF).

Phosphorus

P₄O₁₀

Simple molecular

Acidic

Reacts with water → phosphoric acid (H₃PO₄). Neutralised by bases.

Sulfur

SO₂ / SO₃

Simple molecular

Acidic

SO₂ + H₂O → H₂SO₃. SO₃ + H₂O → H₂SO₄. Neutralised by bases.

Chlorine

Cl₂O₇

Simple molecular

Strongly acidic

Reacts with water → HClO₄. Neutralised by bases.