Chemistry

6.1 - Matter and Atoms

Matter

• a substance that has volume (occupies space) and has a mass

• all matter are made up of atoms which bond together to produce different substances

Atoms

• “building blocks of matter” (classic definition of atoms)

• 1802- first atomic theory of matter presented by john dalton

  • proposed that all matter is made up of tiny spherical particles which are invisible and indestructible

  • now know it is incorrect

  • atoms are made of smaller subatomic particles (protons, neutrons, electrons)

6.2 - The Atomic World

Definitions

Elements: Cannot be broken down into smaller parts by physical or chemical means

Compounds: Made up by two or more elements, can be separated chemically

Molecules: The smallest part of a substance

Compound vs Element: Inseparable vs Separable

Are all elements molecules?: No

Pure substances vs mixtures

• matter can be split into two groups

  • pure substances

  • mixtures

Mixtures

• consists of two or more types of particles that are not chemically combined together

• can be made elements, compounds or both

  • e.g. oxygen gas (O₂) carbon dioxide (CO₂) nitrogen gas (N₂₎

  • e.g. oil-water mixture

• mixtures can be separated by physical means (e.g. filtration)

Pure Substances

• a substance made up of only one type of particle throughout

• fixed composition

• constant properties

• can either be one single element or one single compound

• every sample of substance must contain exactly the same thing

• must have a fixed, definite set of properties

  • pure element - copper metal (Cu)

  • pure compound - (CO₂)

• cannot be separated by physical means (e.g. filtration)

Elements

• made up of just one type of atom

  • could be monoatomic (exist as individual atoms)

• cannot be separated into simpler substances by physical or chemical means

• each element has a unique name and chemical symbol

Compounds

• different types of atoms combine to form new substances

• molecules - two or more atoms that are held together by chemical bonds

Pure Substances: Element or Compound?

• to determine whether a pure substance is an element or a compound, you must determine if the substance can be broken down into simpler substances

How to Identify if Substances are Pure or Mixture

Summary

• elements: cannot be broken down into smaller parts by physical or chemical processes (only one type of atom)

• molecules: two or more atoms that are held together by chemical bonds (which still holds all its properties), smallest part of a substance

• compounds: different types of atoms combined to form a new substance, can be separated chemically

• atoms: the smallest part of a substance that cannot be broken down chemically

• mixtures: two or more types of particles which are not chemically bound together

• pure substances: substances made up of only one type of particle throughout

• the atomic structure (bohr model):

Element VS Compound

Element

• one type of atom

• generally monatomic (only one atom)

• can be di- or poly- atomic, so long as they are the same type of atoms

  • e.g. O, C, H (mono)

  • e.g. H₂, O₂, N₂, C₆₀, (di)

Compound

• two or more atoms chemically bonded (e.g. covalent, ionic bonding)

• these atoms must be different elements

  • e.g. O₂, H₂O, NaCl

Molecule

• two or more atoms that are covalently bonded

• all molecules are compounds but not all compounds are molecules

                    (insert venn diagram)

school of thought 1: any atoms chemically bonded are compounds, any covalently bonded compound is a molecule

• criteria for being a molecule

  • made up of atoms that are covalently bonded

  • made up of atoms of the same or different elements

• criteria for both a molecule and compound

  • made up of atoms that are covalently bonded (makes a molecule)

  • the atoms are made up of two or more different elements (makes a compound)

school of thought 2: any two different chemically bonded atoms are compounds, any covalently bonded same atoms are molecules

6.3 - Atomic Structure

Atomic Structure

• an atom is made up of three sub-atomic particles

  • protons (positively charged)

  • neutrons (neutral charge)

  • electrons (negatively charged)

• note: in chemistry, the word particle is a general term that refers to a small unit of matter. depending on the context, ‘particle’ could mean an atom, a molecule, an ion, or something else

• protons and neutrons are relatively equal mass and are huge in comparison to electrons. this is why the nucleus is dense

• protons and neutrons contribute nearly all the mass of the atom

Rutherford’s Model

• most of the mass of an atom, and all of the positive charge must be located in a tiny central region called the nucleus

• most of the volume of an electron is empty space, occupied only by electrons

• the electrons move in circular orbits around the nucleus

• the force of the attraction between the positive nucleus and the negative electrons is electrostatic

• prior to rutherford’s experiment, the atom was thought to be a spherical cloud filled with electrons all over (think raisin cakes)

• in 1911, rutherford’s experiment proved that the atom had a tiny but heavy nucleus and that most of the volume of an atom is empty space occupied by electrons

The Bohr Model

• in 1913, niels bohr developed a new model of the hydrogen atom that explained emission spectra

• the bohr model proposed the following:

  • electrons revolve around the nucleus in fixed, circular orbits

  • the electrons’ orbits correspond to specific energy levels in the atom

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

• scientists quickly extended bohr’s model of the hydrogen atom to other atoms

• they proposed that electrons were grouped in different energy levels called electron shells. these electron shells are labelled with the number n = 1, 2, 3… 7

  

  • energy levels increase the higher it goes

  • higher the gravitational pull, lower the energy

  • the electrostatic pull of the nucleus slows down the electrons

• two electrons in the same electron shell (energy level) have anticlockwise spins, aka they circle around each other at the same speed meaning that they will never meet

                                    (insert diagram)

• all elements’ electrons orbit in the energy levels at the same speed.

Different Atoms

• the type of atom that makes up each element is determined by the number of protons (atomic number) in the nucleus (protons determine electron identity)

• atomic number: the number of protons in the nucleus of the atom

• mass number: the total number of protons + neutrons in the nucleus

• atoms are electrically neutral: therefore number of electrons = number of protons

  

• (the periodic table shows the most common isotopes)

Calculating the Number of Substance Mass

• example atom: argon

  • the number of protons = Z = 18

  • the number of neutrons = A - Z = 22

  • the number of electrons = Z = 18

• protons determine elemental identity

• electrons determine chemical reactivity

  • (ability of an atom to undergo chemical change in a chemical reaction)

• neutrons determine the physical properties of an element

Isotopes

• all atoms that belong to the same element have the same number of protons in the nucleus and therefore the same atomic number, Z

• atoms that have the same number of protons (atomic number) but different number of neutrons (and therefore different mass numbers) are known as isotopes

  

• isotopes have identical chemical properties but different physical properties such as mass and density

• some isotopes are radioactive

Electron Configuration

Electron Configuration - Bohr Model

• using the bohr model, it is possible to determine the basic electron configuration of any atom by applying the following rules:

  • rule one: each shell can only contain a certain maximum number of electrons

    electron shell	maximum number of electrons
    1	2
    2	8
    3	18
    4	32
    n	2n²

  • rule two: lower energy shells fill before higher energy shells

• these rules only really fully apply to the first 1-18 elements

  • for any element with an atomic number of >18 → shell 3 holds 8, no more 2n² until really heavy elements (start in transition metals)

Worked Example

Flame Test Theory

• energy is the ability to do work and comes in a variety of forms, for example heat, light, motion, electricity, etc

• when atoms are given energy (in this case, heat energy), the electrons jump up to a higher energy level

• work-changing matter

  • chemical state (i.e. in a chemical reaction (process where atoms are rearranged so that bonds are broken and new))

  • movement

  • physical state (liquid, gas, solid

• e.g. kinetic energy is the ability to move matter (work)

  • acceleration

  • deceleration

  • stop moving

  • change direction

• e.g. potential energy is the ability to store energy (work)

• e.g. gravitational energy is the ability to pull objects with smaller mass to objects with bigger mass (work)

• e.g. thermal energy is the ability to combust matter, excite electrons (work)

  • e.g.

    

  • explanation: in n=1, electrons are moving around with 1J (joule), and in n=2, they are moving with 3J of energy, so the electrons need 2J from the flame to move from n=1 to n=2

    • oxygen is used to combust orgainic matter. as long as there is oxygen, any organic matter can combust

Ions

Ions

• remember: protons have a charge of +1, electrons of -1

• a positively or negatively charged atom or group of atoms

  • monoatomic: Na⁺, Cl^-, Cu⁺²

  • polyatomic: Po^3-, OH^-

• charged particles which form when an electrically neutral atom gains or loses electrons

  • this causes to an imbalance, leading to a positively or negatively charged ion

Cations (🐈‍⬛) and Anions (🧅)

• when an atom loses electrons it becomes positively charged, or, a cation

  • e.g. Mg²⁺, Al³⁺

• when an atoms gains electrons it becomes negatively charged, or, an anion

  • e.g. Cl^-, O^2-

• a neutrally charged atom has an equal amount of protons (+) and electrons (-)

  

Octet Rule

• the octet rule is the tendency on atoms to prefer to have 8** electrons in their valence shell*

• this is because this is a stable arrangement, and equivalent to a full shell

• when given the opportunity, atoms will gain, lose or share electrons to follow the octet rule

  • this forms bonds (ionic, covalent, metallic (will be covered later))

* the valence shell is the outermost electron shell

** for hydrogen and helium, the octet rule does not apply, and instead they must have 2 electrons in their valence shell to achieve stability

Cation or Anion?

• atoms always take the path of least resistance

• this means that if, for example, an atom has two in their valence shell, they will loose two atoms, and if they have six, they will gain two atoms

• all metals will become cations, all non-metals will become anions

• if the transition metals are removed, then the remaining groups are numbered 1-8, \ representing the number of electrons in their atoms’ valence shells (excluding He)

• groups 1-3 all shed electrons, while groups 5-7 all gain electrons

• group 8 elements are non-reactive, as they already have a full valence shell

• group 4 metals do not give or take, but instead share, forming covalent bonds (more on this later)

  

Naming Monoatomic Ions

• cations: refer to the metal name

  • e.g. Al³⁺ → Aluminium Cation

• anions: replace the suffix with ‘-ide’

  • e.g. O^2- → Oxide

Electron Transfer and Ionic Bonding

Ionic Compounds

• always formed with a metal and non-metal

  • metallic cation + non-metallic anion = ionic compound

• during a reaction, there is a transfer of electrons (Metal → Non-metal)

• once this occurs, the oppositely charged ionic bond together in a lattice

• ionic bond = electrostatic attraction between a cation and anion

Lattice Structure

• made up of cations and anions side by side

• opposites touching only (✨opposites attract✨)

• e.g. sodium chloride (aka table salt)

  

• no. cations = no. anions

• ionic compounds are brittle, as when a force is applied to the lattice, the ions with like charges align and actively repel each other, shattering the lattice

  

Different Models of Ionic Bonding

Model	Example	Does Not Show
Chemical Formula	NaCl	charges lattice structure
Dot and Cross Diagram		lattice structure ionic bonds
2D Diagram		how ions were formed more than one layer
3D Diagram		charges that there are no spaces between ions

Electron Transfer Diagrams

• used the show the path that electrons take when they are removed from a metal and added to a non-metal during ionic bonding

• e.g.

  

• steps

            1- draw an electron shell diagram of the neutral metal and non-metal

            2- add a ‘+’ between them

            3- draw an arrow leading from each valence electron in the metal to the valence shell

                of the non-metal

            4- add an arrow toward the resulting ions

            5- draw an electron shell diagram of the resulting cation/s and anion/s

            6- write the chemical symbol of the metal and the non-metal in the centre of the

                electron shell diagram, taking care to add charges to your ions

Naming Ionic Compounds

Rules

• name cations (metals) first before anions (non-metals)

  • extension- if the metal is a transition metal, indicate the valency in numerals after the name (e.g. Fe(III), Ag(I), Gold(I))

• the name of the metal cation remains as is (e.g. Sodium ion, Na⁺ ion)

• the name of the non-metal is changed. it’s suffix becomes ‘-ide’

• Na⁺ (cation) + Cl^- (anion) = NaCl (sodium choride)

Rules for Writing Chemical Formulas (e.g. Lithium Oxide)

1. write the symbol and charge of the two ions forming the ionic compound - Li⁺, O²^-

2. calculate the lowest common multiple of the two numbers in the charges of the ions - 1×2=2

3. calculate how many cations and anions are needed to equal the lowest common multiple - two Li⁺ ions, one O²^- ion

4. determine the formula of the ionic compound write the symbols of the cation first - Li₂O

   • the number one should not be written in subscript

shortcut:

            Li^1+|\                O^2-|/

                         Li₂O

Naming Ionic Compound

• the name of ionic compounds are written by listing the name of the positive ion followed by the name of the negative ion. a space separates the two parts of the ionic compound name

  • e.g. NaCl → sodium chloride

  • e.g. K₂O → potassium oxide

  • e.g. CaH₂ → calcium hydride

• if the cation element has more than one possible oxidation state (charge), follow the element name by parenthesis containing the approximate Roman Numerals

  • e.g. Fe₂O₃ → Fe³⁺, O^2- → Iron (III) oxide

• there is no prefix indicating the number of atoms in the cation. so Hg₂Cl₂ is mercury(II) dichloride

* Ca²⁺, O^2- → CaO (simplest form) CaO (wrong)

* x⁴⁺, y^2- → xy² (simplest form) x²y⁴ (wrong)

examples:

• K₂O → Potassium Oxide (K⁺, O^2-)

• NaOH → Sodium Hydroxide (Na⁺, OH^-)

• CaBr₂ → Calcium Bromide (Ca²⁺, B^-)

• Al₂S₃ → Aluminium Sulfide (Al³⁺, S^2-)

• Li₃N → Lithium Nitride (Li⁺, N^3-)

• Be(NO₃)₂ → Beryllium (Be²⁺, NO₃^1-)

Metallic Bonding

Metallic Bonding

• only made up of metals (cations)

  

  • cations shed valence shell electrons

  • delocalised sea of electrons (free moving electrons)

Structure: Crystal Nature of Metallic Bonding

• metals occur as crystal lattices

• this is because metallic atoms tend to lose their outer shell electrons easily

  • this turns the metal elements into positively charged cations

• these cations form a metallic lattice structure, in which electrons from each metal atom overlap with each other, forming a sea of electrons that can flow between all the metal ions.

  

• electrostatic forces of attraction between the positively charged metal cations and negatively charges valence electrons occur in all directions, holding the lattice together

• this type of non-directional bonding is known as metallic bonding

• this means that metal atoms are hard to seperate but relatively easy to move (malleable)

  

Alloys

• alloys are 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 they contain. this is due to the pure metal atoms being the same and arranged in layers, as apposed to alloys which contain elements with atoms of different sizes

  

  • e.g. steel- iron (metal) + carbon (non-metal)

  • e.g. bronze- copper (metal) + tin (metal)

  • e.g. brass- copper (metal) + zinc (metal)

Types of Bonding

• metal + metal = ionic compound

• non-metal + non-metal = covalent compound

• metal + metal = metallic mixture

How to Form a Lattice

• metals → “big granular chunks” (e.g. sodium → big, soft, reactive)

• + non-metals → often a gas (e.g. chlorine → liquid)

• millions of atoms combine at high temp

Metallic vs Ionic Compounds

metals are good conductors of electricity