Atomic Theory Sample exam questions Flashcards
Define atomic orbital
Region of space around the nucleus where electrons are most likely to be found
Define first ionisation energy
The minimum energy required to remove most loosely bound electron from a neutral gaseous element in its ground state
Explain general increase in first ionisation energy across a period
- Increases across
- Decreasing atomic radius
- Increase in effective nuclear charge
- Screening effect remains the same
Define Atomic number
Number of protons in the nucleus of an atom
Define atomic mass
Number of protons and neutrons in the nucleus of an atom
Define Relative atomic mass
The average mass of an atom compared to 1/12th the mass of the carbon 12 isotope
Define electronegativity
The measure of the attraction that an atom has for the shared pair of an electron in a covalent bond
Describe how electronegativity values can be used to predict the type of bonds that occur between a pair of atoms
<0.4= pure covalent bond
>0.4 but <1.7=polar covalent bond
>1.7= ionic bond
explain the trend of atomic radius down the group
- Increases
- nuclear charge increases
- screening effect increases
explain why alkali metals are so reactive
- Group one
- 1 electron in its outer energy level
- it loses that one electron to follow the octet rule
Explain why alkali metals become more reactive down the group
- atomic radius increases down the group
Explain why bromine is less reactive than chlorine
- The electronegativity of Bromine is smaller
- The atomic radius is larger than chlorine
Explain why neon is chemically inert
8 electrons in its outside shell
How is an ionic bond formed between a pair of atoms
- Electrons are transferred to form a positive ion
- To form a negative ion
How is a double covalent bond formed between a pair of atoms
The sharing of 2 pairs of electrons
transition metal
one that forms an ion with an incomplete d-subshell
Give 2 characteristic properties of transition metals
variable valency
acts as a catalyst
what is a crystal
solid with a repeating pattern of atoms
Describe the bonding in a metallic crystal
positive ions with valence electrons being shared
would you expect nickel to be a good electrical conductor, why?
yes
electrons are free to move
Account for the visible lines in the hydrogen emission spectrum
- Electrons occupy the ground state
- they move to excited energy levels due to heat or electricity
- it’s unstable and falls back down
- light is released
Why is there a sharp increase from the 11th to the 12th ionisation energy
12th electron is taken from a new energy level which is closer to the positive nucleus
Intermolecular forces
forces that exist between molecules
Name the type of intermolecular force between
i) ammonia gas
ii) water
iii) methane gas
i) hydrogen bonds
ii) hydrogen bonds
iii) van der Waals forces
explain in terms of intermolecular forces why the NH3 is very soluble in water
- water and ammonia are both polar molecules
- one side of each is slightly negative and the other side is positive
- hydrogen bonds can form between them and ammonia dissolves
- “like dissolves like”
Relative atomic mass
The average mass of one element compared to 1/12th the carbon 12 isotope
Why do electronegative values decrease across a period
- nuclear charge increases
- electrons enter the same energy level
- screening effect stays the same
- Electronegativity values decrease as the atomic radius increases
Covalent crystal bonding
Each carbon atom forms four covalent bonds with four more carbon atoms
What are the bonding forces in an ionic crystal
Ionic bonds
What are bonding forces that hold molecular crystals together
Intermolecular forces
Explain why the molecular crystals usually have lower melting points than ionic crystals
Intermolecular bonds in molecular crystals are weaker than ionic bonds which are easier to break
What type of crystal is characterised by a solid having a lustre and very high mobility of electrons throughout the structure
Metallic
Atomic Number
Number of protons in an atom of a nucleus
First ionisation energy
The minimum energy required to remove the most loosely bound electron from a neutral gaseous element in its ground state
A general increase in ionisation energy
- Nuclear charge is increasing
- Atomic radius is decreasing
- Screening effect stays the same
why is there a lot of energy required to ionise a neon atom
Neon is a noble gas
It has 8 electrons in its outer shell
It is stable and does not want to lose an electron
Why is the first ionisation energy of xenon higher than neon
Xenon has a larger atomic radius than neon
Define atomic orbital
A region in space around the nucleus where electrons are most likely to be found
What’s the difference between the 2s and 2p orbitals of neon
S is lower in energy than the P orbital
Px, Py, Pz orbitals differ in direction
Why do electron pairs repel each other
Like charges repel
Electrons are negatively charged so they repel each other at an energy level
Compare the magnitude of repulsion between the two lone pairs and repulsion between two bonding pairs of electrons
Repulsion between two lone pairs of electrons is greater than the repulsion between two bond pairs
Lone pairs of electrons are closer to the nucleus than bond pairs
Distinguish between ground and excited state
Ground: electrons occupy the lowest available energy level
Excited: Electrons occupy a higher energy level
2 Ways an electron can become excited
Heat
Electricity
Why is there no yellow line in the hydrogen emission spectrum
No transition equal in energy to produce a yellow line
Electrons cant occupy spaces between energy levels
Describe how you carried out a flame test
- Soak a wooden splint in water
- dip in salt
- place splint into the blue flame of a bunsen burner
Define mass number
Number of neutrons and protons in an atom of a nucleus
Define relative atomic mass
Mass of an atom of an element compared to 1/12th the carbon-12 isotope
Define electronegativity
The measure of the force of attraction of an atom in a molecule for the shared pair of electrons in a covalent bond
Identify a type of bond in a water molecule
Polar covalent bond
Predict solubility of methane
Methane is not soluble in water
water is polar
methane is non-polar
Like dissolves like
