A1 Atomic Structure + Bonding Flashcards
Define an isotope
- Atoms of the same element
- With different numbers of neutrons and different masses
Define Relative Atomic Mass
- The weighted mean mass of an atom of an element
- Compared to 1/12th the mass of an atom of Carbon-12
Define Relative Isotopic Mass
- The mass of an atom of an isotope
- Compared to 1/12th the mass of an atom of Carbon-12
Define Weighted Mean Mass
- The mean mass
- Taking into account the relative abundancies of the isotopes
What does the atomic number of an atom tell you
Number of protons
What does the mass number of an atom tell you
The number of neutrons and protons
How do you calculate the RFM given abundances and relative atomic masses
On a mass spectrometer, what do the peaks indicate
Isotopes
What happens to all elements when in a mass spectrometer
Develop a 1+ charge
How to calculate the RFM from a mass spectrometer, where the relative abundances don’t equal 100
What peaks should be shown on mass spectrometers with diatomic elements
All possible combinations of diatomic molecules, and all monotomic atoms
How many electrons occupy the first shell
2
How many electrons occupy the 2nd shell
8
How many electrons occupy the 3rd shell
18
How many electrons occupy the 4th shell
32
How many orbitals in s subshell
1
How many orbitals in p subshell
3
How many orbitals in d subshell
5
How many electrons in ANY orbital
2
Define an orbital
- A 3D reigon of space around the nucleus, that can hold up to 2 electrons
- With opposite spins
What is the shape of an s orbital
What is the shape of a p orbital
What is the exception to the rule of subshells filling with increasing energy
4s fills before 3d
How do electrons fill sub-shell orbitals as you go along the period
Singularly, before doubling up in each orbital
Why is an element in the x block
As it’s highest energy electrons occupies the x subshell
When writing electronic configurations, how should you write it when the 3d subshell is full
Write it before 4s
What electrons are removed when positive ions form from d block elements
4s are removed before 3d
Electronic configuration for chronium
Electronic configuration for copper
Define Ionic Bonding
- The strong electrostatic attraction
- Between positive and negative ions
Define an Ionic Lattice
A repeating pattern of oppositely charged ions
When drawing ionic lattices, what must be included in each circle
Both the formula and the charge
Why do ionic compounds have high melting/boiling points
- Giant ionic lattice structure
- Has lots of very strong ionic bonds
- Lots of energy is required to break all the strong ionic bonds
What ions are always soluble
- Na+
- K+
- NH4+
- NO3-
State and explain the conductivity of solid ionic compounds
- Does not conduct
- Ions are fixed in position
- Not mobile and free to move
- Around giant ionic grid-like lattice structure
State and explain the conductivity of aqueous and molten ionic compounds
- Does conduct
- Ions aren’t fixed in position
- Ions are mobile and free to move
- Around the giant ionic grid-like lattice stucture
Define covalent bonding
- The strong electrostatic attraction
- Between a shared pair of electrons and the nuclei of bonded atoms
Define a Dative Covalent Bond
A covalent bond where both electrons are donated by one atom
What is used to measure the strength of a covalent bond
Average bond enthalpy
Why does diamond/silicon have a very high melting/boiling point
- Each carbon atom makes 4 strong covalent bonds with other carbon/silicon atoms
- Lots of energy is needed to break all the strong covalent bonds
- In the giant tetrahedral covalent lattice
Why does graphite/graphene have a very high melting/boiling point
- Each carbon atom makes 3 strong covalent bonds with other carbon atoms
- Lots of energy is needed to break all the strong covalent bonds
- In the giant covalent lattice structure
State and explain the conductivity of diamond/silicon
- Cannot conduct
- Each carbon/silicon atom makes 4 strong covalent bonds with other carbon/silicon atoms
- There are no delocalised electrons, or ions mobile and free to move
- Around the giant tetrahedral covalent lattice strucutre
State and explain the conductivity of graphite/graphene
- Does conduct
- Each carbon/silicon atoms makes three strong covalent bonds with other carbon atoms
- There is one delocalised electron per atom mobile and free to move
- Around the giant covalent lattice structure
Why is graphite soft
- Layers of carbon atoms are held together by weak forces of attraction
- Layers are able to slide over each other
Which giant covalent structures are soluble
None of them
Why do simple molecular substances have a low boiling point
- Molecules held together by weak intermolecular forces
- Not a lot of energy is needed to overcome the weak intermolecular forces that act between molecules
State and explain the conductivity of simple molecular substances
- None conduct
- As all the molecules are neutral
- As theres no delocalised electrons or ions present
Define Metallic Bonding
- Strong electrostatic attraction
- Between a lattice of cations and a sea of negatively charged delocalised electrons
Draw a metallic lattice
Why do metallic compounds have high melting/boiling points
- There are lots of strong metallic bonds
- In the giant metallic lattice structure
- Lots of energy is required to break all the strong metallic bonds
State and explain the conductivity of metals in all states
- Conducts in all states
- As the sea of delocalised electrons are mobile and free to move
- Around the giant metallic lattice structure
When are metallic substances soluble
Never
Draw a linear molecule
Draw a trigonal planar molecule
Draw a tetrahedral molecule
Draw a trigonal bipyramid molecule
Draw a octahedral molecule
Draw a trigonal pyrimid molecule (NH3)
Draw a bent molecule (H2O)
What are the relative repulsions of lone electron pairs and bonded electron pairs
Lone pairs repel more than bonded pairs
How to write an answer, comparing bond angles
- State the number of bonding reigons and lone electron pairs
- Lone pairs repel more strongly than bonded pairs
- Bonded electron pairs repel equally
Define electronegativity
- The ability of an atom to attract the bonding electrons
- In a covalent bond
What is the most electronegative element
Flourine
How to determine what element is more electronegative
Whatever is closest to flourine in the periodic table
In a covalent bond, which atom do the electrons move closer to
The more electronegative atom
How to explain if this molecule is or isn’t polar
- Has polar bonds
- Molecule is symmetrical
- Dipoles cancel out
- Non-polar
How to explain if this molecule is or isn’t polar
- Has polar bonds
- Molecule isn’t symmetrical
- Dipoles don’t cancel out
- Polar
Where do all intermolecular forces act
Between molecules
In what molecules do London forces occur
In every molecule (NOT giant covalent)
Why as molecules get larger, do their boiling points increase
- More electrons
- Stronger London forces
- More energy required to break London forces
How are London forces induced
- Electrons move randomly in molecule
- Creates a temporary dipole in the molecule
- Induces temporary dipoles in neighbouring molecules
Describe a permanent dipole force
- Polar molecules have dipoles
- Dipoles interact to form the dipole-dipole force
Between what molecules does hydrogen bonding occur
- One lone pair of flourine, oxygen or nitrogen
- Hydrogen of another molecule
What is the relative strength of the different intermolecular forces
- Hydrogen bonding strongest
- Permenant dipole-dipole forces
- London forces
Draw a hydrogen bonding diagram for two water molecules
What must be included in all hydrogen bonding diagrams
- Lone pairs
- Dipoles
- Hydrogen bond
Describe and explain the anomolous properties of ice - relativly high melting point
- Hydrogen bonding is very strong
- So lots of energy needed to overcome it
Describe and explain the anomolous properties of ice - ice is less dense than water
- Water molecules held apart in an open lattice structure
- By hydrogen bonds
State and explain the solubility of non-polar substances in non-polar solvents
- Soluble
- IMF’s form between molecules in solvent and molecules in solute
- Weakens the IMF’s in the simple molecular solvent
State and explain the solubility of polar substances in non-polar solvents
- Insoluble
- Attraction between molecules in solvent and ions in ionic lattice
- Not strong enough to break ionic bonds in ionic lattice
State and explain the solubility of polar substances in polar solvents
- Soluble
- Polar bonds in solute
- Attract polar bonds in solvent
