VCE Chemistry Unit 1 Notes

Elements and the Periodic Table

• Music and chemistry are both universal languages.

• Elements and the Periodic Table: VCE Chemistry Unit 1 AOS 1.

• Topic: How do the chemical structures of materials explain their properties and reactions?

Atomic Structure

• Element: Pure substance containing only a single type of atom.

• Atom: Smallest unit of matter used in VCE chemistry.

• Nucleus: Region at the center of an atom that contains protons and neutrons.

• Protons (+) are attracted to the electrons (–) in an atom.

• Proton: Positively charged particle inside the nucleus of an atom.

• Neutron: Neutral particle inside the nucleus of an atom.

• Electron: Negatively charged particle that exists outside the nucleus of an atom.

• Subatomic particle: Particles that exist inside the atom.

Atomic Number

• Each element has its own atomic number.

• The atomic number represents the number of protons in the nucleus of an atom.

• The current periodic table is ordered by atomic number.

• Atomic number: The number of protons in the nucleus of an atom.

• Chemical symbol: An abbreviation used to represent a chemical element.

• Periodic table: Table of chemical elements in which elements are arranged in order of increasing atomic number.

• Atoms have the same number of protons and electrons because they have no overall charge.

• When an atom loses or gains electrons, it becomes an ion which is a charged particle.

• Ion: An atom that has lost or gained electrons to become a charged particle.

Mass Number

• Protons and neutrons have a similar mass.

• Electrons have negligible mass, so their mass is not considered when doing calculations.

• The mass number of an element is a relative value compared to other elements and therefore does not have any units.

• Mass\ number = no.\ protons + no.\ neutrons

• Mass number: Mass of an atom, approximately equal to the average sum of protons and neutrons.

• Number\ of\ neutrons\ in\ an\ atom = mass\ number – atomic\ number

Isotopes

• All atoms of carbon contain 6 protons.

• Isotopes of carbon: ^{12}C, ^{13}C, and ^{14}C.

• Atoms of the same element can have different mass numbers.

• Isotope: Variants of an element which have the same atomic number but a different number of neutrons in their nuclei.

• Isotopic notation (Scientific notation).

• Isotopic notation is on the VCE periodic table.

Molecules and Compounds

• When two or more atoms bond by sharing electrons, they form molecules.

  • Example: H_2 hydrogen gas

• When the atoms are from different elements, then a compound is formed.

  • Example: NH_3 ammonia

• Molecule: When two or more atoms bond by sharing electrons.

• Compound: Two or more atoms of different elements bonded together.

The Periodic Table (Part 1)

• The periodic table is an organizational tool to identify patterns and trends.

• Topic: How do the chemical structures of materials explain their properties and reactions?

Mapping the Periods

• Periods: Rows in the periodic table.

Mapping the Groups

• Groups: Columns in the periodic table.

• Noble gases: Elements in Group 18.

• Elements are placed in order of their atomic number, which represents the number of protons in each element.

• The majority of the 118 elements in the periodic table are metals.

• Metals are shown in blue.

• Non-metals are located towards the right-side of the table shown in green.

• Metalloids (semi metals) are the ‘staircase’ shown in red.

• Metals: Type of elements found toward left hand side of periodic table.

• Non-metal: Type of elements found toward right hand side of periodic table.

• Elements in the same group (column) have similar reactivity.

• Elements in the same period have the same number of occupied electron shells.

• Group 1 (Alkali Metals) and Group 17 (Halogens) are very reactive.

• Group 18 (Noble gases) are generally inert.

Electron Shells

• Electrons are found in shells.

• There are 2 electrons in the energy level closest to the nucleus.

• The number of electrons in each energy level at ground state.

• The number listed first is closest to the nucleus.

• Electron configuration: Arrangement of electrons in shells and/or subshells.

• Valence electrons: Electrons in the outermost shell of an atom.

• Valence shell: Outermost energy shell where valence electrons are found.

• Energy shells: Orbits containing different levels of energy, around the nucleus.

• Electron shell notation: Summary of the number of electrons per shell every level in the format a, b, c, d. Commas separate shells.

• Ground state: When electrons are at their lowest possible energy level.

  • Distance from the nucleus. Shell number: max number of electrons in the shell

  • closest 1 2

  • 2 8

  • 3 18

  • farthest 4 32

Blocks

• Elements belong in blocks according to the location of their valence electrons.

• s–block: Groups 1 and 2.

• p–block: Groups 13 to 18.

• d–block: Groups 3 to 12.

• f–block: Lanthanoids (atomic number 57 – 71) and actinoids (atomic number 89 – 103).

• Blocks: Assortment of elements into discrete categories based on which subshell their valence electrons are found in.

Subshells and Orbitals

• The Schrodinger model includes orbitals within each electron shell

• The Pauli Exclusion Principle states that each orbital can hold 0, 1, or 2 electrons only.

• Orbitals: Regions with the highest probability of finding electrons.

• Pauli Exclusion Principle: Rule that states an orbital can’t hold more than 2 electrons.

• Electron subshell notation: Order of filling of sub-shells including s,p,d,f.

Electron Subshell Notation

• The electron sub-shell notation gives the number of electrons in each energy sub-shell.

  • e.g. 1s22s22p63s1

  • The 1st number is the shell number.

  • The lower-case letter is type of sub shell.

  • The superscript number equates to number of electrons in a particular sub-shell.

  • Exceptions for copper and chromium.

• Aufbau principle: Rule that states that subshells are filled by electrons from the lowest to the highest energy level.

  • s sub-shell fills first

  • Collect shell number together

  • 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f

Core Charge

• Core charge is the effective pull of the positive nucleus on the valence electrons.

• Core charge increases across a period.

• Core charge is constant in a group.

• Core\ charge = number\ of\ protons – inner\ shell\ electrons

• Core charge: The number of protons in the nucleus minus the number of inner-shell electrons.

Atomic Radius

• Core charge and the number of electron shells affect atomic radius.

• Decreases from left to right across a Period to Group 18.

• Increases going down a group.

• Atomic radius: The distance from the center of an atom to the valence electrons.

• Atomic radii: Plural of atomic radius.

The Periodic Table (Part 2)

• Properties are affected by periodicity.

• Periodicity: Characteristics of elements in a period.

• Going across a period, the characteristics of elements change.

  • increasing core charge across a period

  • decreasing atomic radius across a period

  • same number of occupied electron shells

Electronegativity

• Electronegativity increases across a period and up a group.

• Electronegativity increases with:

  • increasing core charge

  • decreasing atomic radius

  • fewer occupied electron shells

• The greater the electronegativity, the greater the ability of the atom to attract electrons towards its nucleus.

• Fluorine is the most electronegative element.

• Caesium and francium have the lowest electronegativity.

• Noble gases aren’t considered to have electronegativity because they have a full valence shell of electrons.

• Electronegativity is very important in bonding.

• The quantitative measurement of electronegativity is called the Pauling Scale.

• One way of remembering the order of electronegativity is FONCl, or FOClN.

• Electronegativity: The ability of an element to attract shared electrons towards itself.

First Ionization Energy

• An ion is an atom that has lost or gained electrons.

• Elements vary in the amount of energy required to remove the first valence electron.

• First ionization energy increases across a period as core charge increases and atomic radius decreases.

• First ionization energy decreases going down a group as the number of occupied electron shells increases and the valence electrons are further from the nucleus.

• First ionization energy: The energy required to remove the first valence electron from an atom.

Metallic Character

• Metals in groups 1 and 2 possess high metallic character. Includes luster, silvery appearance, and the ability to conduct electricity.

• Metallic character depends on how easily electrons are lost.

• Non-metals possess high non-metallic character

• Metalloids display a mixture of metallic and non-metallic properties.

• Metallic character: Degree to which an element is shiny and conducts electricity.

Reactivity

• Reactivity describes how easy it is for an atom to lose or gain electrons

• Noble gases (Group 18) are inert

• Amongst the metals Group 1 are the most reactive followed by Group 2 metals.

• Amongst the non-metals Group 17 are the most reactive, followed by Group 16 non-metals.

• Reactivity: The tendency of an atom to lose or gain electrons.

Recycling Critical Elements

• Topic: How do the chemical structures of materials explain their properties and reactions?

• There are many elements that are in limited supply, called critical elements.

• Elements shaded red are estimated to have 5-50 years until known reserves are depleted.

• Elements shaded orange are estimated to have 50–100 years until known reserves are depleted.

• Elements shaded yellow are estimated to have 100–500 years until known reserves are depleted.

• This periodic table version was published in 2013.

• Stocks of helium on Earth (and also uranium) are actually becoming depleted.

• Help Protect PRAM

  • Help – helium

  • Protect – phosphorus

  • P – Post transition elements

  • R – Rare earth elements

  • A – And

  • M – Metalloids

• Critical elements: Elements that are in short supply and unless recycled, sources could become depleted in 50–100 years.

Helium

• 2nd most abundant element in the universe (after hydrogen).

• Less dense than air – can escape Earth’s atmosphere.

• Inert – useful in a non-reactive atmosphere.

• Used in the production of semiconductors, fibre-optic wire and LCD panels.

• Used in liquid form as a coolant in MRI machines and nuclear reactors.

Phosphorus

• Key element in living organisms – including in bones, DNA, ATP, and cell membranes.

• Used as fertilizer.

• Too much phosphorus from fertilizer can cause algal blooms in waterways called eutrophication.

• Food waste contains phosphorus.

• We need to consider recycling phosphorus from contaminated water ways.

Rare Earth Elements

• Refers to 17 metallic elements including lanthanoids, plus scandium and yttrium.

Post-Transition Elements

• Groups 13-16 are post-transition elements.

• Example – gallium

  • Has a low melting point – useful application in thermometers

  • Used in semiconductors and computer chips

  • Present in LEDs

• Post-transition elements: Elements that are located between the transition metals and metalloids on the period table.

Linear Economy vs Circular Economy

• In order to be able to use critical elements, we need to move from a linear economy to a circular economy.

• Four of the 12 green chemistry principles are related to a circular economy.

  • Atom economy

  • Design for energy efficiency

  • Prevention of wastes

  • Use of renewable feedstocks

• Sustainable development: Meets the needs of the present without compromising the ability of future generations to meet their own needs.

• Linear economy: Operates on a ‘take-make-dispose’ model, making use of resources to produce products that will be discarded after use.

• Circular economy: A continuous cycle that focuses on the optimal use and re-use of resources from the extraction of raw materials through to production of new materials, followed by consumption and re-purposing of unused and waste materials.

• Green chemistry principles: Principles aimed at reducing the chemical-related impact on both humans and the environment through dedicated sustainability management programs.

Covalent Bonding (Part 1)

• Covalent bonding involves the sharing of electrons.

• Topic: How do the chemical structures of materials explain their properties and reactions?

• Types of Bonding Between Atoms:

  1. Covalent

  2. Ionic

  3. Metallic

• Covalent bonding involves the sharing of electrons in:

  • diatomic molecules

  • larger molecules containing three or more atoms

• Ionic bonding is where electrons are transferred between atoms, and a metal ion and non-metal ions are attracted to each other.

• Metallic bonding is where two or more metal atoms are bonded.

• Covalent bonding: The sharing of electrons between non-metal atoms.

• Molecule: Two or more non-metal atoms chemically bonded together.

• Diatomic molecule: Two non-metal atoms covalently bonded, of same or different elements.

• Covalent bonding occurs when non-metal atoms share electrons in an intramolecular bond in order to satisfy the octet rule.

  • C has four valence electrons.

  • N has three valence electrons.

  • O has two valence electrons.

  • H and halogens have one valence electron.

• Octet rule: General principle stating that atoms are more stable if they have 8 electrons in their valence shell.

• Intramolecular bond: Covalent bond within a molecule.

• Single covalent bond: Sharing of one pair of electrons.

• Double covalent bond: Sharing of two pairs of electrons.

• Triple covalent bond: Sharing of three pairs of electrons.

• One pair of electrons shared between two non-metal atoms is known as a single covalent bond.

Formulas

• The molecular formula can be seen as a summary of the number and type of atoms in a molecule.

  • Does not show orientation of atoms, nor shape of molecule.

  • Ex: serotonin C10H12N2O

  • Molecular formula: The number and type of atoms in a molecule.

• Electron dot structures show the valence electrons.

  • For less than four valence electrons, they are shown distributed around the element symbol.

  • For more than five valence electrons, they are represented in pairs.

  • A pair of valence electrons is called a lone pair and is not involved in covalent bonding.

  • Electron dot (Lewis) structure: A representation showing valence electrons as dots.

  • Lone pair: A pair of valence electrons not involved in covalent bonding.

• In the structural formula, lines represent a pair of shared electrons

  • Double lines between atoms indicate a double covalent bond.

  • Triple lines between atoms indicate a triple covalent bond.

  • Structural formula: Representation of a molecule using lines to represent a pair of shared electrons.

• Skeletal formula: Representation of a molecule where vertices are C atoms and H are assumed to be added so C has 4 bonds.

• Space–filling model: Representation of a molecule demonstrating the three-dimensional arrangement and the respective shape and size of atoms.

• Ball and stick model: Representation of a molecule demonstrating the three-dimensional arrangement.

Covalent Bonding (Part 2)

• The shape of molecules depends on the repulsion of electron pairs according to valence shell electron pair repulsion (VSEPR) theory.

• Polar and non-polar character with reference to the shape of the molecule.

• Lone pair: A pair of valence electrons not involved in covalent bonding.

• VSEPR theory: Model used to predict the different shapes of molecules based on the extent to which electron pairs repel each other.

Molecular Shapes

• Tetrahedral: a 3D geometric shape with four triangular faces. Each bond is at an angle of 109.5°.

• Pyramidal: a 3D geometric shape with three faces.

  • Each bond is at an angle of 107°.

• Bent (V–shaped): Each bond is at an angle of 104.5°.

• Linear: Each bond is at an angle of 180°.

MMolecule Polarity

• Both the shape of a molecule and the types of atoms involved determine the polarity of a molecule.

• Non-metals differ in their electronegativity.

• We can use the Pauling scale of electronegativity to determine the difference in electronegativity, and therefore the polarity of bonds.

  • Less than 0.4: Non-polar covalent bond

  • Between 0.4 and 2.0: Polar covalent bond

  • More than 2.0: Ionic bond

• If a molecule contains only non-polar bonds, then the molecule will be overall non-polar.

• If a molecule is perfectly symmetrical and contains polar bonds, then the molecule is overall non-polar.

• If a molecule contains at least one polar bond, and the molecule is asymmetric then it is a polar molecule.

• The molecule is called a dipole.

• A molecule that contains a polar covalent bond is not necessarily polar overall.

• Polarity: Property of having a partial positive and negative charge.

• Electronegativity: How strongly an atom attracts bonding electrons towards itself.

• Polar covalent bond: Covalent bond between atoms with an unequal distribution of electrons in the bond.

• Non-polar covalent bond: Covalent bond between atoms where electrons are shared equally.

• Permanent dipole moment: Covalent molecules with a permanent positive and negative charge due to the difference in electronegativity between the atoms.

• Polar molecule: Molecule with a partially positively charged end and a partially negatively charged end.

Intramolecular Bonding and Intermolecular Forces (Part 1)

• Intramolecular bonding is very strong covalent bonding, much stronger than intermolecular forces.

• Show covalent bonds as lines.

• Types of intermolecular forces include:

  • dispersion forces

  • dipole-dipole attraction

  • and hydrogen bonding

• Covalent bonding: The sharing of electrons between atoms.

• Intramolecular bonding: Very strong covalent bonds between atoms in molecules.

• Intermolecular forces: Forces between molecules. Weaker than intramolecular bonding.

• Intra means within, inter means between e.g. intranet, internet.

• All molecules are attracted to each other by weak intermolecular forces called dispersion forces.

• They are the weakest type of intermolecular force.

• Dispersion forces:

  • Occur because electrons are constantly moving.

    • When electrons spontaneously move to the far left or far right of an atom, an instantaneous dipole moment is created.

    • A temporary negative side of a molecule is attracted to a temporary positive side of another molecule.

• As the number of electrons in atoms and molecules increases, so does the strength of the dispersion forces.

• Intermolecular forces are weaker than covalent bonds, so are shown as dotted lines.

• Dispersion forces: Weakest type of intermolecular forces due to instantaneous dipoles between adjacent molecules.

• Dipole-dipole attraction is stronger than dispersion forces between molecules.

• This is due to permanent dipole-dipole moments instead of temporary dipole-dipole moments.

• The strongest type of dipole-dipole attraction is hydrogen bonding.

• Dipole-dipole attraction: Electrostatic force of attraction between a partial positive charge on one molecule with a partial negative charge on another molecule.

• Occurs when there is a significant difference in electronegativity between atoms of different molecules.

• Hydrogen bonding:

  • It is the strongest form of both intermolecular and dipole-dipole attraction.

  • Occurs between H on one molecule bound to F or O or N and F, O or N on another molecule.

  • Due to the presence of lone pairs of electrons on F, O, or N.

  • Hydrogen bonding: Strongest form of intermolecular forces. Occurs between a H bonded to F, O or N on one molecule and an adjacent F or O or N atom on a neighbouring molecule.
    FonCl or FOCIN 2B:30

Intramolecular Bonding and Intermolecular Forces (Part 2)

• The properties of molecular substances are determined by their intermolecular forces.

• Intermolecular forces are responsible for:

  • melting point

  • boiling point

  • hardness
    The strength of intermolecular forces reflects the amount of energy it takes to break the intermolecular forces.

• It takes more energy to break hydrogen bonds than dipole-dipole and dispersion forces.

• Gases have more kinetic energy than liquids.

• Liquids have more kinetic energy than solids.

• Kinetic energy: The energy of an object due to its motion.

Melting Point

• Solids have lower kinetic energy, and therefore have stronger intermolecular forces between molecules than the same substance in a liquid state.

• When there is enough energy to overcome the intermolecular forces the substance can change state into a liquid.

• The melting point temperature depends on the type of intermolecular forces present.

• Melting point: The temperature at which a substance changes its state from a solid to a liquid.

Boiling Point

• Gases have higher kinetic energy than liquids.

• Liquids have intermolecular forces between adjacent molecules.

• As molecular size increases, so too does the number of dispersion forces between adjacent molecules.

• Therefore, as molecular size increases so too does the melting point and boiling point.

• The more linear a molecule, the more adjacent molecules that can pack together.

• Therefore, the more opportunity for dispersion forces between linear molecules, leading to higher melting and boiling points.

• Harder substances have greater intermolecular forces between adjacent molecules, than less hard substances.

• Boiling point: The temperature at which a substance boils and changes its state from a liquid to a gas.

• Isomers: Molecules with the same molecular formula.

• Hardness: Resistance to deformation when subjected to pressure.

Macromolecules

• Carbon is a very versatile element.

Carbon

• Carbon is a non-metal.

• Carbon atoms have 4 valence electrons.

• Carbon has three naturally occurring isotopes 12C, 13C and 14C.

• Carbon is the building block of life due to its ability to form complex, stable molecules.

• Each allotrope of carbon has its own structure.

• Allotrope: One of the different physical forms of an element due to different structural arrangement of atoms.

• Analogy – graphite and waffles.

• Lattice: Regular three dimensional arrangement of atoms.

• Network: Atoms bonded in a continuing arrangement extending outwards.

• Covalent network lattice: Regular three dimensional arrangement of atoms covalently bonded together extending outwards.

• Covalent layer lattice: Atoms covalently bonded together in a two dimensional arrangement to form layers, held together by weak dispersion forces.

Diamond

• High melting point

• Very hard

• Brittle

• Does not conduct electricity

• High thermal conductivity

• Insoluble

• Hardness resists wear and enhances durability in cutting tools

• Strong covalent bonding enables thermal conductivity in electrical components

• Diamond lasers can be produced as optical components

• Able to induce friction on other objects without being worn itself as an abrasive

Graphite

• High melting point

• High thermal conductivity

• Soft, slippery feeling

• Less dense than diamond

• Insoluble

• Able to conduct electricity, can transfer current from a stationary wire to moving parts in electric motorsCarbon brushes in electric motors

• Inert, electrically conductive material in Electrodes in batteries

• Layers of graphite can slide over each other, reduces friction in machinery for industrial lubricant

Metals

• Metals have low ionization energies

• Therefore, relatively small amounts of energy is required to remove valence electrons from atoms. When metal atoms lose one or more valence electrons, they form cations

• ore: deposit in Earth's outermost layer containing metals and other minerals

• metallic bonding: the electrostatic force of attraction between delocalized electrons and cations in a metallic lattice structure

• Sea of delocalized electrons: electrons that move freely between metal cations in the metallic bonding model

• Crystal lattice: atoms fo one type of metal element that are metallically bonded and organized in a pattern

  • metallic bonding is non-directional as the electrostatic forces of attraction between the cations and delocalized electrons are in all directions (unlike ionic and covalent bonding). Sea of delocalized electrons model Sea of delocalized electrons

  • Metallic bonding is strong because of the electrostatic force of attraction between the delocalized electrons and the cation lattice

Malleability

* metals can be bent and hammered into different shapes without breaking
* Malleable:    *Ductile: ability to be hammered or stretched into a thjn shape without breaking
* Electrical conductivity:  *HeatConductivity
* lusture: shiny and glossy appearance
* high melting and boiling point: can withstand high temperatures, without melting or boiling

Metallic Bonding Model

g## MetalReacticity
reactions

Metal WaterReactions. 270

Hydroxice ion (oH-)
General reactions for Reactive metal and water
water +reactive metal => metal hydroxide + Hydrogen
Groups 1 ,2
MISTAKEN GENERALIZATION. 276
REACTIONS OF IONIC COMPOUNDS: VSEPR THEORY
ionic bond: icalium
Anious
solutions of ionic will swapped
* The other ionic compound of solids. => precipitate. then solution
* C-O-B_A-L-T colou, odour, bubble, appearance of colid light/ sound, temperature change
* solubility tale can used