Mod 1 Chemistry

investigate the role of electronegativity in determining the ionic or covalent nature of bonds between atoms 🖳

What is electronegativity?

Electronegativity is a measure of an element's ability to attract/ pull electrons when atoms of the element share electrons with another element

Pauling Scale

investigate the differences between ionic and covalent compounds through:

1. using nomenclature, valency and chemical formulae (including Lewis dot diagrams)

NOMENCLATURE

In ionic bonding:

-Compounds end  in -ide

-The cation is first with unchanged name

-Roman numerals used to show which ion with transition metals (oxidation state)

In covalent bonding:

-prefix changes depending on number of atoms

-exceptions are halides and organic compounds

-more electronegative element named last

CHEMICAL FORMULAE (INCLUDING LEWIS STRUCTURE)

In ionic bonding:

-Represents a ratio, therefore can be simplified

-Within Lewis structure, ions are bracketed and charge assigned to each ion.

-Big number to show number of ions, dots indicate anion valence shell

In covalent bonding:

-Subscript indicates number of atoms per element

-Within Lewis structure, atoms should satisfy octet rule

-shared pairs can be represented with a stem

-Dots used to represent valence electrons of each element

VALENCY

-Metals always form positive ions, non-metals always (except hydrogen) form negative ions

-In ionic compounds, electrons are exchanged/donated

-In covalent, electrons are shared

1. examining the spectrum of bonds between atoms with varying degrees of polarity with respect to their constituent elements’ positions on the periodic table

2. modelling the shapes of molecular substances

-Only covalent compounds can form molecules

-Non-polar bonds -> equally shared electrons

-Polar bonds -> unequally shared electrons

Vectors and delta indicate polarity:

Different Molecular shapes:

1. Linear:

1. Bent

1. Trigonal planar

1. Pyramidal

1. Tetrahedral

*Both shape and electronegativity determine polarity

investigate elements that possess the physical property of allotropy

What are allotropes?

Allotropes are different physical forms of the same element, possessing the same chemical properties but different physical traits (boiling and melting point, conductivity, hardness…)

*Remember: Carbon, Oxygen and phosphorus

investigate the different chemical structures of atoms and elements, including but not

limited to:

1. ionic networks

Can be shown with:

Empirical formula (simplified ratio, least detail)

Examples:

1. covalent networks (including diamond and silicon dioxide)

• Shared electrons in 3D lattice

• Empirical formula (Because infinite lattice)

• Insoluble in water

• No molecules formed

1. covalent molecular

• Molecular formula

• When intermolecular forces loosened in solid, it melts

• When intermolecular forces broken in liquid, it evaporates

• Poor conductors of heat

• Covalent substances are not soluble in water unless polar!!! (because they attract water molecules through hydrogen bonding)

1. metallic structure

• Insoluble in water

• Usually strongest chemical bond that arises from electrostatic attractive force between electrons

explore the similarities and differences between the nature of intermolecular and intramolecular bonds and the strength of the forces associated with each, in order to explain the:

• physical properties of elements

• physical properties of compounds

What are the types and nature of intermolecular forces?

Dispersion forces:

• Between all molecules

• Affected by size, shape and number of electrons

• Weakest intermolecular forces

• Temporary

• Dipole-induced attraction, London forces, Van Der Waal's forces

What affects the strength of these forces?

More electrons + bigger atom = more distance of which electrons can move (also valence electrons are less tightly held) = higher probability of temporary dipoles (and stronger interactions when temporary dipoles)= bigger dispersion forces

Greater molecular mass = Higher dispersion forces (therefore more energy needed to loosen or break them) = higher MP and BP

Dipole-dipole forces

• Between molecules of a substance that is polar

• Stronger than dispersion

• Attraction between partially positive end of polar molecule and partially negative end of other molecule

• Strength decreases with size because polarity decreases as poles are further apart

Hydrogen Bonding

• Strongest intermolecular force

• Involves hydrogen with the very electronegative elements:

1. Fluorine  2. Oxygen   3. Nitrogen

Examples:

H₂O, NH₃, HF

What are the types of intramolecular forces?

1. Ionic bonding

2. Covalent bonding (between atoms)

3. Metallic bonding

How do they affect properties of elements and compounds?

Atomic structure and atomic mass

IQ2: Why are atoms of elements different from one another?

Students:

investigate the basic structure of stable and unstable isotopes by examining:

•  their position in the periodic table

Elements that are further down on periodic table are have larger atomic mass and therefore, are more unstable

•  the distribution of electrons, protons and neutrons in the atom

In stable isotopes, the amount of protons is the same or more than the amount of neutrons.

When there are more neutrons than protons in a nucleus, the isotope is unstable

Number of electrons is unchanged in isotopes

•  representation of the symbol, atomic number and mass number (nucleon number)

model the atom’s discrete energy levels, including electronic configuration and spdf notation

*discrete energy levels means energy shells

What are the different ways electron configuration can be represented?

2n² to find max number of electrons in each shell

1. Condensed

Example: Argon = 2,8,8

1. Orbital notation/SPDF notation

-Based on Schrodinger model where electrons occupy 3D space around nucleus called 'orbital'

-Assumes electrons have wave-like properties and therefore, have shells AND subshells AND orbitals

-Aufbau principle describes filling order where lowest energy orbitals are always filled first

-Pauli Exclusion principle states that each orbital can contain max of 2 electrons, with each electron having different spin

1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²4d¹⁰5p⁶6s²4f¹⁴5d¹⁰6p⁶7s²5f¹⁴6d¹⁰

What does large number represent? shell number

What does the letter represent? Orbital/subshell

What does superscript represent? Number of electrons

*Exceptions: Chromium 2,8,12,2 -> 2,8,13,1 and Copper 2,8,17,2 -> 2,8,18,1

1. Filling diagram

-Hund's Rule: Every orbital in a subshell is singly occupied with one electron before any orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin

Example:

calculate the relative atomic mass from isotopic composition

What is relative atomic mass?

The relative atomic mass of an element is the average mass of the atoms present in the naturally occurring element relative to the mass of an atom of the carbon-12 isotope taken as exactly 12

Calculate:

-Average atomic mass = (mass of the isotope x relative abundance)

e.g.

Average atomic mass of carbon = 12 x 0.99 + 13 x 0.01 = 12.01 amu

-Relative atomic mass from isotopic composition-

A_r =

-To calculate percentage abundance of each isotope from the mass spectrum:

Measure peak height with ruler then find total (by adding peak heights) to find percentage abundance

investigate energy levels in atoms and ions through:

• collecting primary data from a flame test using different ionic solutions of metals

Results:

Is it the cations or the anions that produce distinctive colour?

Anions in nitrate produce orange colour, and anions in chloride produce orange pinkish colour.

Expected results:

Why are there differences in results?

The colour of the flame may have changed observation. The water may cool down the flame, reducing energy and producing different colour

•  examining spectral evidence for the Bohr model and introducing the Schrödinger model

What is an emission spectra?

An emission spectra consists of a series of bright coloured lines against a black background. It is produced by a gas when heated or electricity makes  it get so hot that it emits light of particular wavelength

Why do we see a particular colour?

• When atoms get excited, they move to higher energy levels then fall back down to 'ground state' after a short time

• The excess energy is released as light. The greater the energy change, the higher the frequency of the emitted photon

• Electron returning to ground state can return in a number of ways

• However, Bohr's model could not quantitatively predict interpret spectra more complex than a hydrogen atom.

ENTER SCHRODINGER MODEL:

• Treated electrons as wave instead of particle

• Schrodinger equation was successfully used to interpret the emission spectra of atoms with many electrons

investigate the properties of unstable isotopes using natural and human-made radioisotopes as examples, including but not limited to:

•  types of radiation

•  types of balanced nuclear reactions

  

Properties of Matter

IQ1: How do the properties of substances help us to classify and separate them?

Students:

explore homogeneous mixtures and heterogeneous mixtures through practical investigations:

•  using separation techniques based on physical properties

Homogeneous

1. Evaporation

2. Distillation

3. Chromatography

Heterogenous

1. Magnetism

2. Sieving

3. Separating funnel

4. Filtration

•  calculating percentage composition by weight of component elements and/or compounds

Recall gravimetric analysis

Percentage composition =

Percentage yield =

investigate the nomenclature of inorganic substances using International Union of Pure and Applied Chemistry (IUPAC) naming conventions

What is an inorganic substance?

A substance which does not contain carbon bonded with hydrogen

3 types of inorganic compounds:

Type 1: Group I and II metal ->  cation anion-ide

Type 2: Transition metal -> cation (roman numeral) anion-ide
Type 3: Non-metal -> prefix-nonmetal prefix-nonmetal

classify the elements based on their properties and position in the periodic table through their:

•  physical properties

•  chemical properties

demonstrate, explain and predict the relationships in the observable trends in the physical and chemical properties of elements in periods and groups in the periodic table, including but not limited to:

•  state of matter at room temperature

•  electronic configurations and atomic radii

• Group number = number of electrons in outer shell

Atomic radii:

How can you account for this trend?

• Each Period adds a shell which increases distance of electrons from nucleus. This also decreases nuclear charge so attractive force is reduced with each cell, increasing radius

• Less protons compared to electrons means weaker pull on electrons so bigger radius

•  first ionisation energy and electronegativity

What is ionisation energy?

The energy needed to remove an electron from outer shell (when the element is in gaseous state)

How can you explain this trend?

As number of protons increases, pull on electrons is stronger (nuclear charge). As group increases, nuclear charge decreases with added electron shell, meaning weaker pull on electrons and less energy required to take them from atom. Also atoms in groups 1-3 want to give away electrons to reach stable composition

Electronegativity:

How can you account for this trend?

Increases towards the right because higher nuclear energy means stronger pull on electrons

Decreases down period because added shells reduce pull of electrons by nucleus (therefore affected by atomic size)

•  reactivity with water

• Group I are very reactive and all react with cold water to produce hydrogen and a base (metal hydroxide)

• Group II are reactive as well, reacting with water (except magnesium which reacts with steam) with same products above

• Metal oxides react with water to force bases

• Non-metal oxides react with water to form acids

• Amphoteric oxides can react with both acids and bases