periodic table

how did mendeleev arrange elements in the periodic table?

• in order of atomic mass

• left gaps

• predicted properties of undiscovered elements which would go into gaps

• similar properties

how does the modern periodic table arrange elements?

• by increasing atomic number

what is the periodic table arranged into?

groups and periods

what is the periodic table split into?

s, d, and p blocks

define metaloids and what group classes as metaloids ?

properties of both metals and non- metals

what are group 7 elements also known as?

halogens

what are group 0 elements also known as?

noble gasses

what are group 1 elements also known as?

alkali metals

define periodicity?

repeating trends in properties of the elements across each period in the periodic table

what are ionization energies a measure of?

a measure of how easily an atom loses electrons to form positive ions

define first ionization energy?

• the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions

is ionisation endothermic or exothermic and why?

endothermic as it takes in energy

3 important points about ionisation energy?

• must use gas state symbol

• always refer to 1 mole of atoms

• the lower the ionization energy, the easier it is to form an ion

• the higher the ionization energy, the stronger the force of attraction between the electrons and the nucleus

does ionization energy increase or decrease down a group

decreases down a group, easier to loose outer electrons

what factors affect ionization energy?

• nuclear charge

• atomic radius

• electron shielding

why do ionization energies decrease as you move down the group?

• as you move down the group elements further down have more electron shells compared to those above

• the extra shells mean there is a larger atomic radius

• this reduces their attraction to the nucleus

• the extra inner shells shield the outer electrons from the attraction of the nucleus

why do ionization energies increase across a period?

• gets harder to move outer electrons

• because the number of protons is increasing

• as the positive charge of the nucleus increases, electrons are pulled closer to the nucleus decreasing atomic radius

when do successive ionization energies occur?

• each time an electron is removed

state the 3 giant covalent lattices?

• diamond

• graphene

• graphite

define giant covalent lattices?

huge networks of covalently bonded atoms

why can carbon from covalent structures?

because carbon can form 4 strong covalent bonds

which of the 4 covalent lattices is the hardest?

diamond

in what shape are atoms arranged in diamond?

tetrahedral shape

properties of diamond?

• high melting point

• hard

• cant conduct electricity

• cant dissolve in any solvent

properties of graphite?

properties of graphene?

• a single layer of hexagonally arranged carbon atoms

• delocalised electrons free to move along the sheet

• conducts electricity

• extremely strong

• transparent and light

what structures do metal elements exist as?

giant metallic lattice structure

properties of metals?

• high melting point

• high boiling points

• malleable as metal ions are free to slide over each other

• good thermal conductors as delocalised electrons can pass kinetic energy to each other

• good electrical conductors as delocalised electrons can move and carry charge

• ductile

• insoluble

how can metals conduct electricity?

• delocalised electrons move freely throughout the metal lattice carrying charge therefore allowing them to conduct electricity

how are metals malleable?

• due to positive ions being able to slide over each other

comparison between covalent bonds and metallic bonds?

• covalent bonds are localised whereas metallic bonds are delocalised

• covalent bonds are shared pairs of electrons

what do more atoms in a molecule result in?

stronger induced dipole di[ole forces

why do both melting and boiling points increase across period 2 and 3?

metallic bonds get stronger as the ionic radius decreases and the number of delocalised electrons increases

why do simple molecular structures have weak melting and boiling points?

because they have weak intermolecular forces

which noble gases have the lowest melting and boiling points

neon and argon as they are held by the weakest forces

what are group 2 metals also known as?

alkaline earth metals

does ionisation energies increase or decrease down group 2?

decreases due to increasing atomic radius and sheilding

what do group 2 elements lose electrons to form?

they loose electrons to form positive ions

what do group 2 elements react with water to produce?

hydroxides and hydrogen

what do group 2 elements react with oxgen to produce?

oxides

what do group 2 elements react with dilute acid to produce?

salt and hydrogen

what do the oxides of group 2 elements react with water to form?

metal hydroxides, the hydroxide ions make solutions strongly alkaline

where is calcium hydroxide used and what is it used to neutralise

used in agriculture to neutralise acidic soils

what are magnesium hydroxide and calcium carbonate used to treat?

used in indigestion tablets as antacids

what are group 7 elements also known as?

halogens

name the first 4 halogens?

• fluorine

• chlorine

• bromine

• iodine

what molecules do halogens exist as?

diatomic molecules

define diatomic molecules?

2 atoms joined by a single covalent bonds

why do boiling and melting points increase down group 7?

• due to increasing strength of london forces and the increase of relative mass of atoms

what is the physical state and colour of fluorine?

gas, pale yellow

what is the physical state and colour of chlorine?

gas, green

what is the physical state and colour of bromine?

liquid, red-brown

what is the physical state and colour of iodine?

solid, grey

how and what do halogens react together to form?

by gaining electrons to form 1- ions

are halogens oxidised or reduced?

reduced

why are larger atoms less reactive/ why does reactivity decrease down the group?

• atomic radii increases as you go down the group

• outer electrons get further away from the nucleus

• harder for larger atoms to attract outer electron needed to form an ion

• so larger atoms are less reactive

when will a halogen displace a hallide?

if a halide is below it in the periodic table