Periodic Table Flashcards
what is in S-block
alkali metals
alkali earth metals
what is in D-block
transition metals
what is in P-block
halogens and noble gasses (right side of table)
what is in F-block
lanthanides and actinides
what does period and groups show?
period- energy levels
group- amount of valence electrons
Properties of metals?(5)
- shiny solids (lustrous)
- good conductors of heat and electricity
- malleable(bend into shapes) and ductile (stretched into wire)
- high melting and boiling points
- react with acids and give off hydrogen gas
Properties of semi-metals?(5)
- solids at room temp
- non malleable/ ductile
- resists flow of electricity
- when chemical reaction with metals, act like non metal
- when chemical reaction with non metals, act like metal
Properties of non-metals?(5)
- dull
- insulators of heat and electricity
- some are gasses at room temp, rest are solids
- gaseous non metals have low melting and boiling points
- do not react with acids
how do elements change moving from left to right in periodic table?
metallic nature decreases
how do elements change moving down groups of periodic table
densities and melting points change
how do metals change moving down group
become more reactive
how do non metals change moving down group
reactivity decrease
define ionisation energy?
energy needed per mole to remove an electron from an atom in the gaseous phase
define First ionisation energy?
energy needed per mole to remove the first electron from an atom in the gaseous phase
Explain why energy increases to remove successive electrons from an atom? (2)
atom becomes more positively charged with electrons being removed. more stronger attractive force between nucleus and electrons. more closer the electrons to the nucleus the more energy required to remove them
Why does noble gasses require most energy to remove electron?
Noble gasses already have full outer shells and are STABLE making it harder to remove electron thus requiring high energy.
why so group one and 2 have low ionisation energies?
easily form positive ions to be stable and thus less energy is needed to remove electrons.
what does high ionisation of group 15- 17 show?
unlikely to give away electrons when they react and are more likely to gan electrons than lose them during chemical reactions to be stable.
difference between stable and neutral?
stable- full outer shell- low energy
neutral- same amounts of positive and negative charges
trends of ionisation energy? (2)
first ionisation energy decreases down a group
first ionisation energy increases from left to right across a period.
why does ionisation energy increases from left to right across a period?
increasing atomic number means greater effective nuclear charge which means there is a stronger attractive force to the nucleus meaning more energy is needed to remove an electron.
explain 2 reasons why ionisation energy decreases down a group?
- valance electrons are further away from nucleus meaning weaker force of attraction to nucleus meaning less energy is needed to remove valance electron
- more core electrons as you move down group. more core electrons repel valance electrons and shielding effect will increase. this results in a decreased effective nuclear charge causing valance electrons to experience weaker force of attraction and thus, less energy needed to remove electron
what happens when ionisation occurs in the next energy level of an atom?
the electron needs to be removed from the next level meaning it is closer to nucleus and more closer the electrons to the nucleus the more energy required to remove them as stronger attractive forces.
define electron affinity
amount of energy released per mole when an atom or molecule gains an electron to form a negative ion.
trends of electron affinity? (2)
electron affinity increases across a period and decreases down a group
why does electron affinity decrease down a group?
higher in the group, the less shielding effect there will be as there is less energy levels meaning elements higher in the group will release more energy when forming ions.
what happens if an element needs to gain an electron to become stable?
it will release more energy when gaining electron
Define electronegativity
tendency of an atom to attract a shared pair of electrons to itself
whats is electro negativity measured in?
Pauling scale
why does electro negativity increase across a period?
each time there is one more proton meaning stronger nuclear attraction on a bonding pair of electrons
why do noble gasses not have electro negativities?
do not form bonds
What is atomic radius?
measure of size of an atom
what is ionic radius?
measure of the size of a cation or anion
why does radius decrease across period?
more protons= greater effective nuclear charge= stronger force of attraction pulling energy shells closer together.
why does radius increase down group?
more core electrons as you move down a group and shielding effect
Radius of cation?
smaller radii of atoms from which formed as there is greater nuclear charge attracting electrons to nucleus
Radius of anion?
larger radii that atom from which formed as increased shielding effect= greater repulsion
