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UNIT 2: CHEMISTRY

1) Evolution of the Atom:

Matter - A substance that has volume (occupies space) and mass (weight). All matter is made up of atoms, which bond together to produce different substances.

Atoms - The building blocks of mater (classic definition)

First atomic theory: In 1802, John Dalton proposed that all matter is made up of tiny spherical particles that were indivisible and indestructible. We now know that this is wrong.

2) The Atomic World

Pure substances and mixtures:

Matter can be split into two types; pure substances and mixtures

Pure substances - Substances made up of only one type of particle throughout. They cannot be physically separated.

Mixtures - Substances that are made up of two or more types of particles throughout. These can be physically separated.

Elements:

Elements are types of atoms that share the same number of protons. In nature, some elements prefer to be formed in a monatomic form (individual atoms), while others form in a polyatomic form (e.g. molecules)

Regardless of the preferred arrangement, elements can also be considered pure substances, because if you have multiple particles, they will all be the same.

Compounds:

Compounds are substances that are made up of two or more different elements that are chemically bonded together. If they are not mixed with anything else, these compounds can also be considered pure substances, but not all pure substances are compounds.

3) Atomic Structure and Isotopes:

Atomic structure:

An atom is made up of three different types of particles:

• Protons - positively charged

• Neutrons - neutrally charged/no charge

• Electrons - negatively charged

The protons and neutrons make up the nucleus, which has majority of the mass of an atom, due to its dense center.

The Bohr model:

In 1913, Niels Bohr developed a new model of the hydrogen atom, which explains emission spectra.

The Bohr model proposed that:

• Electrons revolve around the nucleus in fixed, circular orbits

• The electrons’ orbits refer to specific energy levels in atoms

• Electrons can only occupy fixed energy levels and cannot exist between two energy levels

Scientists soon evolved the Bohr model to atoms other than hydrogen. They proposed that electrons were grouped in different energy levels, called electron shells. They will labelled with the number n = 1, 2, 3, etc.

Different types of atoms:

The type of atom that makes up each elements are determined by the number of protons (atomic number) in the nucleus.

Atomic number = number of protons

Mass number = number of protons + number of neutrons

Atoms are electrically neutral - number of electrons = number of protons

Isotopes:

Isotopes are when an atom has the same number of protons, but different number of neutrons. They have identical chemical properties, but different physical properties.

4) Electron Configuration:

Electron configuration - Bohr model:

Using the Bohr Model, the basic electron configuration can be found for different types of atoms. Remember, this rule ONLY applies to the first 18 elements.

Shell 1 - 2 electrons

Shell 2 - 8 electrons

Shell 3 - 18 electrons

An overall approach can be 2, 8, 8, …

There are two rules to the electron configuration model:

Rule 1 - Each shell can only contain a specific number of electrons

Rule 2 - Lower energy shells must be filled before higher energy shells

6) Ions:

The octet rule:

The octet rule is the tendency of atoms to prefer having eight valence electrons, because a valence shell with eight electrons is very stable, and equivalent to a full shell. When given the opportunity, atoms will either gain, lose or share electrons to follow the octet rule.

Ions:

Ions are atoms with a charge. They are formed when a neutral atom gains or loses electrons. Losing electrons can cause an imbalance between positive and negative charges, thus creating an ion.

When an atom gains electrons, it is more negative. If it loses electrons, it becomes more positive.

There are two types of ions; cations and anions.

Cations - positively charged ions

Anions - Negatively charged ions

To become stable, ions take the path of least resistance, or the most effective way, between either losing or gaining valence electrons.

For example, if an ion has 3 valence electrons, it will lose all of them to be stable.

Metals and nonmetals:

Metals have low amounts of valence electrons (1, 2, 3), so they usually lose electrons and become cations.

Nonmetals have high amounts of valence electrons (5, 6, 7) and they gain electrons and become anions, with the exception of noble gases (group 18), since they already have eight valence electrons.

Naming monatomic ions:

With cations, you put the word ‘cation’ next to the element.

e.g. Aluminium cation

With anions, replace the suffix with -ide

e.g. Oxide anion

7) Electron Transfer and Ionic Bonding:

Ionic bonding:

Ionic bonding is the electrostatic attraction between two oppositely charged atoms. They typically involve a metal cation and a nonmetal anion.

Lattice structures:

Some ionic compounds involve more than two ions. As more ions bond together, they form an ionic lattice.

Electron transfer diagrams:

Electron transfer diagrams are used to show what path electrons take during ionic bonding.

NOTE - for lattices, cations are always next to anions, and vice versa.

8) Naming Ionic Compounds:

Rules for naming ionic compounds:

• Names of cations (metals) should come before names of anions

  e.g. sodium chloride

• Names of cations remain constant

  e.g. sodium

• Names of anions have -ide at the end

  e.g. chloride

Chemical formulas for simple ionic compounds:

Rules for writing chemical formulas:

• Write the symbol of the positively charged ion first

• Use subscripts to indicate the number of each ion, writing them after the ion they match to.

• If there is only one ion of an element, don’t put subscript 1.

• Don’t include the charges of ions in the balanced formula.

Example:

In this example, the Lithium symbol is first, since it is positively charged. The Oxygen ion is negatively charged, so it is second. The number of the negative charge of the Oxygen ion goes as a subscript next to the Lithium ion, and vice versa.

More examples:

K₂O = K¹⁺O^2- Potassium Oxide

NaOH = Na¹⁺OH^1- = Sodium Hydroxide

CaBr₂ = Ca²⁺Br^1- Calcium Bromide

Al₂S₃ = Al³⁺S^2- = Aluminium Sulfide

Li₃N = Li¹⁺N^3- = Lithium Nitride

Be(NO₃) = Be²⁺NO₃- = Beryllium Nitrate

9) Properties of Ionic Compounds:

Ionic compounds are called ‘salts’, and they are solid in room temperature. They have strong electrostatic bonds between cations and anions and form a lattice. Very high levels of energy are required to break these bonds, so they stay solids.

Why are ionic compounds brittle?

When an ionic compound is struck with a hammer, the ions shift, and go next to like ions. Due to this, they repel, and the bond is broken.

Some ionic compounds are soluble, since the ions are attracted to water’s particle charges, and therefore, the bonds break. Since the atoms are free, they can conduct electricity, hence, solid ionic compounds can’t conduct electricity, but liquids can.

Molten ionic compounds can conduct electricity, but solids can’t, due to the kinetic/thermal energy of the liquid compounds loosens the bonds between the ions, and making them conductive, due to the free moving ions.

10) Metallic Bonds:

Ionic bonds:

• Valence electrons from the cation go to the anion

• The cation (transition metal)‘donates’ its valence electrons to the anion (nonmetal).

• Solid crystals with repeating patterns of cations and anions.

Covalent bonds:

• They share valence electrons

• Usually liquid or solid

• Between two nonmetals

Metallic bonds:

• Metallic bonds are surrounded by electron clouds (delocalised sea of electrons)

• Valence electrons move freely around the metal ions.

• There is only one type of metal cation

• They are uniformly structured

• Occurs in pure metals

• Occur as ‘crystal lattice”.

A crystal lattice of a metallic bond. The red circles are cations, while the blue circles are the delocalised sea of electrons.

This type of non-directional bond is called metallic bonding.

Metal atoms are hard to separate, but relatively easy to move (malleable).

Alloys: A mixture of two or more elements, where one element is a metal combined via metallic bonding.

Alloys are generally harder than the pure elements, due to the cations having different sizes and radii.

Examples of alloys:

Steel: combination of iron (metal) and carbon (nonmetal)

Bronze: Combination of copper (metal) and tin (metal)

Brass: Mixture of copper (metal) and zinc (metal)

Metallic properties:

Metals are:

• Lustrous (due to the presence of free moving electrons, which can reflect light.

• Good conductors of electricity (due to the delocalised sea of electrons being free to move, and an electric current being able to form)

• Malleable (can be hammered into shapes without breaking)

• Have a high melting point (due to strong metallic bonding)

• Good conductors of heat (delocalised sea of electrons are able to gain kinetic energy, then collide with other electrons and conduct heat)

• Ductile (able to be drawn into a thin wire without breaking)

• Dense (has a tightly packed lattice)

Flame test logbook task:

The colour of the flame changes, due to the energy levels of the metallic cations. The heat makes their electrons jump to higher energy levels. As these excited electrons jump back down to their original energy levels, they release the energy in the form of light. The colour depends on the amount of energy released. Each metal has a unique outcome.

11) Covalent Bonding:

Covalent bonds refer to bonds that involve sharing electrons, instead of donating, to satisfy the octet rule, and always involves two or more non-metals.

There is only one pair of shared electrons, called a ‘bonding pair’.

One bonding pair is called a single bond, two bonding pairs are a double bond, etc.

All halogens can covalently bond with themselves, or other halogens.

Steps to draw covalent bonds:

• Draw the symbol of the element that will form the most bonds in the center

• Draw the other atoms around the central atom

• Determine how many electrons are needed to achieve chemical stability, either following the octet rule or the duet rule

• The atom which needs the most atoms to achieve chemical stability is the one that needs to be in the middle

• Add a pair of electrons between any two atoms that are forming a covalent bond

• Determine how many electrons each atom has left and add these as non-bonding pairs

• Check to see that the octet or duet rule is satisfied for all atoms and that all the valence electrons have been drawn.

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UNIT 3: PHYSICS

1) Introduction to Electricity:

Electricity:

Electricity is the movement of energy through charged particles, such as electrons or ions.

Electric charge can be either positive or negative, opposite charges attract and similar charges repel.

Electrons:

Electrons are the most involved in electricity.

Most materials have electrons tightly bound to their atoms, but with electrical conductors (like metals), the electron (delocalised sea of electrons) are free to move.

Conductors and Insulators:

Conductors are able to conduct, because of their electrons that are free to move.

Insulators, however, have electros that hold on tightly to their atoms, and electrons don’t move along these materials very well.

Static electricity:

Some insulators can get electrons from conductors when rubbed together.

The insulator has a more negative charge, while the conductor has a more positive charge.

Static electricity is the abundance of electric charge on a surface.

Static electricity can only be generated by two materials rub against each other, an electric charge is present, and there is an electrical imbalance.

When two materials rub against each other, there are loose electrons, that escape to the other object, making one positively charged, and the other negatively charged.

Like magnets, opposite charges attract, while like charges repel.

Static electricity practical investigation:

Hypothesis:

The object that loses electrons will be positive, and the object that the positive object sticks to will be negative.

Results:

Questions:

1. How did you determine the charge on each material after they had been in contact?

   The material that the other material was stuck to had a negative charge, due to the loose electrons. The object that is attracted has a positive charge, due to loss of electrons.

2. Illustrate this charge differential with a picture.