FA1 Exam Flashcards
What are the three subatomic particles and their charges and mass
Protons: Positive +1, Mass=1
Neutrons: Neutral 0, Mass=1
Electrons: Negative -1, Mass 1/2000th of pro
What are ions and what are they called
Atoms that gain or lose electrons
Gains and electron - anion, negatively charged
Loses and electron - cation, positively charged
What are energy levels, what are the designated letters
Energy levels or shells are regions around the nucleus in which electrons move in. Subshells are within these designated by s, p, d and f.
Afbau principle
Electrons fill orbitals in order of increasing energy
Hund’s rule
Electrons must occupy all orbitals singly before pairing up
Pauli exclusion principle
no two electrons in an atom can have the same set of quantum numbers.
What are the exclusions to the subshells and why
Chromium (Cr) and copper (Cu) exhibit exceptional electronic configurations because half-filled and fully-filled d orbitals are more stable than partially filled ones, leading to a redistribution of electrons between the 3d and 4s orbitals.
Isotopes
Isotopes are atoms of the same element with different number of neutrons, affecting their mass but not their chemical properties. They have the same number of electrons, these electrons determine how these atoms behave in a reaction.
How to calculate relative atomic mass
(rim x ra) + (rim x ra) = relative atomic mass
Relative isotopic abundance x relative abundance for each isotope.
Atomic radius
The atomic radius is the distance from an atom’s nucleus to the boundary of its electron cloud. These distances are small and are measured in picometers.
How does the atomic radius look on the periodic table
- It increases from right to left on the periodic table as the number of protons decrease, minimising their positive charges. The nucleus becomes less positive causing the electrons to be less attracted to it.
- It increases from top to bottom as more electrons increase the number of shells, increasing the radius.
Valencies
Number of electrons in the valence shell, can be calculated for the elements in groups 1, 2 and 13-18 as it is the last digit of the groups number.
Ionic radius
The loss or gain of electrons imbalances the atoms charge, changing the attraction forces within the ion, impacting the ionic radius. It increases further down the group you go as the number of electron shells increase.
First ionisation energy
The first ionisation energy is the amount of energy required to remove one electron from an atom in its gaseous form.
How does First ionisation energy look on the periodic table
Down a Group: Decreases as additional electron shells weaken the attraction, making outer electrons easier to remove.
Across a Period: Increases due to more protons strengthening attraction, making electrons harder to remove.
Exceptions occur due to electron configurations.
Electronegativity
Electronegativity is an atom’s ability to attract electrons towards itself when combined with another atom. It increases from left to right and upwards along a group.
An increased number of inner electrons reduces the attraction of the outer electrons to the nucleus.
Noble gas bonding?
Noble gases are monoatomic, they exist as a single atom. They don’t bond with other atoms due to their stable electron configuration.
Non-noble gas bonding?
Non-noble gas atoms have to bond with other atoms to fill their valence shells. Forming bonds can involve losing, gaining or sharing electrons.
Three types of bonding
Metallic
Ionic
Covalent (network and molecular)
What bonds involve ions
Metallic and ionic. Atoms are already neutral as they have the same number of electrons and protons balancing out each other, however the change in electrons effects this.
Metallic bonding
Metal atoms form a lattice structure, surrounded by each other.
Valence electrons are attracted to multiple nuclei and move freely between atoms.
These delocalised electrons aren’t tied to a single atom and are shared by many, giving metals their unique properties.
Metallic lattice
Metal atoms lose electrons, becoming positively charged cations.
Cations form a lattice structure, surrounded by a “sea” of delocalised electrons.
The electrostatic force between cations and delocalised electrons holds the lattice together.
Metallic bonding properties
- Generally high melting and boiling points
- High electrical conductivity due to the mobility of electrons
- High thermal conductivity, as free electrons transport heat efficiently
- Strong and a generally ductile and malleable
- Hardness depends on metal
- Generally insoluble in most solutes, due to strong electrostatic attraction.
Ionic bonding
Ionic bonds form between metal and non-metal ions. Formed between oppositely charged ions, because the electrons lost and gained must be equal to create a neutral substance.
How to name ionic compounds
- The metal is named first then the non-metal
- The non-metal ends with ‘ide’
Eg. AlCl3 = aluminum chloride
Transition metals
Transition metals have different electron structures to other metals and can show more than one possible charge. e.g. Fe 2+, iron(ii)
Ionic bonding properties
- Melt and boil at very high temps, lots of energy is required to break strong bonds
- Can only conduct electricity good when molten or dissolved in water
- Moderate thermal conductivity, less efficient than metals
- Strong bond strength, but are brittle, and can shatter/fracture
- Typically hard and rigid due to strong ionic bonds
- Solubility depends on strength of attraction between ions and strength of bonds
Covalent bonding
Formed between non-metals. Covalent compounds do not accept of donate electrons, they share electrons to fill their valence shells. The electrostatic force of attraction between the positive nuclei and the shared electrons holds the atoms together.
What are the electrons called in covalent bonds
- The electrons shared between atoms are bonding electrons of bonding pairs
- Electrons not shared are non-bonding pairs or lone electrons
How are covalent bonds called
First element closest to the left of the periodic table has the same name, and the end of the second is changed to ‘ide’. Same group name lower one first.
Exception - when bonded to halogen (cl, Br, I), oxygen is name last.
If more than one element include prefix
What are the prefixes
Mono, Di, Tri, Tetra, Penta, Hexa, Hepta, Octa, Nona, Deca
Molecular covalent bonding properties
- Electrons shared are held together by strong covalent bonds
- Low melting and boiling points due to weak intermolecular forces
- Do not or generally poor electrical conductivity due to localised electrons
- Poor thermal conductors due to weak intermolecular forces
- Much weaker than network, relying on weak molecular forces to hold together in s solid state.
- Many molecular covalent substances are relatively soft.
Network covalent bonding properties
- High melting and boiling points
- One of the strongest materials in the world.
- Extremely hard
Polarity
In a covalent bond, the electrons are not shared equally. A greater electron density is held around the ore electronegative element. When the difference in electronegativities of a compound results in an asymmetrical (uneven) distributions of charges, the molecule is polar.
Nonpolar
When the electronegativity of a compound results in a symmetrical distribution of charges, the molecule in nonpolar.
Polar solubality
A polar solvent such as water can dissolve many polar or ionic solutes because its molecules have the same type of strong forces acting among them as do the potential solutes.
Polar covalent compounds are often soluble in water, where nonpolar ones are not soluble, only in oil and fats.
