3..., 13... Flashcards
(65 cards)
1
Q
period
A
- row of elements (horizontal)
- number of shells
2
Q
group
A
- column of elements (vertical)
- number of valance electrons
3
Q
s,d,p,f blocks on periodic table
A
4
Q
what do all s-block elements have?
A
- only s electrons in the outer shell
5
Q
what do all p-block elements have?
A
- atleast one p-electron in the outer shell
6
Q
what do all d-block elements have?
A
- atleast one d-electron and one s-electron but no f or p electrons in thr outer shell
7
Q
what do all f-block elements have?
A
- atleast one f-electron and atleast one s-electron, but no d or p electrons in the outer shell
8
Q
what is the pattern seen in chemical and physical properties on the periodic table
A
- physical and chemical properties of elements in the periodic table show clear patterns related to the position of each element in the table
- elements in the same group show similar properties
- properties change gradually as you go across a period
9
Q
what is atomic radius?
A
- size of an atom
- ## distance between the nucleus of an atom and the outermost electron shell
10
Q
atomic radius across a period?
A
decreases
11
Q
atomic radius down the group?
A
increases
12
Q
ionic radius
A
- measure of the size of an ion
13
Q
ionic radius down the group?
A
increases
14
Q
ionic radius across a period
A
- ionic radius increases with increasing negativ charge
- ionic radius decreases with increasing positive charge
15
Q
first ionisation energy
A
- ## amount od energy required to remove one mole of electrons from one mole of atoms of an elemts in gaseous state to form one mole of gaseous ions
16
Q
first ionisation energy across a period?
A
increases
17
Q
first ionisation energy down the group
A
decreases
18
Q
Why does ionisation energy increases across a period?
A
- nuclear charge increases
- distance between the nucleus and outer electron remains reasonably constant
- shielding by inner electrons remains the same
- rapid decrease in ionisation energy between the last element in one period and the first element in the next period
- increased distance between the nucleus and outer electrons
- increased shielding by inner electron
- these outweigh nuclear charge
19
Q
why does ionisation energy decrease down the group?
A
- distance between the nucleus and outer electron increases
- shielding by inner shell electrons increases
- effective nuclear charge is decreasing as shielding increases
20
Q
why does successive ionisation energie of an element increases?
A
- ## removing an electron from a positive ion is more difficult than from a neutral atom
21
Q
electron affinity
A
- ## amount of energy released when one mole of electrons is gained by one mole of atoms of an element in the gaseous state to form one mole of gaseous ions
22
Q
electron affinity down the group?
A
decreases
23
Q
electron affinity across a period?
A
increases
24
Q
Electronegativity
A
- ## ability of an atom to attract a pair of electrons towards itself in a covalent bond
25
electronegativity down the group?
decreases
26
electronegativit across a period
increases
27
how does nuclear charge affect electronegativity?
- attractions exist between protons and electrons
- increase in the number of protons leads to an increase in nuclear attraction
- increased nuclear charge/protons = increased electronegativity
-
28
how does atomic radius affect electronegativity?
- electrons closer tot he nucleus are more strongly attracted towards it's positive nucleus
- increased atomic radius = decreased electronegativity
29
bonding of metals and non-metals
metals: metallic due to loss of outer shell electrons
non-metals: covalent by sharing outer shell electrons
30
electron conductivity of metals vs non-metals
metals: good conductors of electricity
non-metals: poor conductors fo electricity
31
type of oxides metals vs non-metals
metals: basic oxides (few amphoteric)
non-metals: acidic oxides (some neutral_
32
reactions with acids metals vs non-metals
metals: many reaction with acids
non-metals: do not react with acids
33
physical characteristics metals vs non-metals
metals: malleable, can be bent and shaped
non-metals: flaky and brittle
34
melting and boiling point metals vs non-metals
metals: high melting and boiling point
non-metals: low melting and boiling point
35
pH of oxides across a period
basic --> amphoteric --> acidic
Al: amphoteric
36
melting point in period three oxides
Na2O: high
MgO: High
Al2O3: very high
SiO2: very high
P4O10: low
SO2: low
37
chemical bonding of oxides across period 3
Na2O: ionic
MgO: ionic
Al2O3: ionic
SiO2: covalent
P4O10: covalent
SO2: covalent
38
Na2O reaction with water and the pH
Na2O+H2O --> 2NaOH
pH:14
39
MgO reaction with water
MgO + H2O --> Mg(OH)2
pH:10
40
P4O10 reaction with water
P4O10 + 6H2O --> 4H3PO4
pH:2
41
NO2 reaction with water
2NO2 + H2O --> HNO3 + HNO2
pH:1
42
SO2/SO3 reaction with water
SO2 + H2O --> H2SO3
SO3 + H2O --> H2SO4
pH:1
43
patter of oxides reacting with water
- metallic oxides form hydroxides when reacted with water
- non-metallic oxides form oxoacids when they react with water
44
how to predict oxides reacting with water
- position of an element in the periodic table can be used to predict and explian its metallic and non-metallic behavior
- metal and non-metal elements generarly form ionic compounds
- oxides become more ionic as you go down the group
- oxides become less ionic across a period
-
45
physical properties of the group 1 metals
- soft eady to cut, softer and denser as you go down the group
- shiny silvery surfaces
- conducct heat and electricity
- low melting and densities
- melting point decreases down the group
46
chemical properties of grp 1 metals
- react readily with oxygen and water vapour in air, so they are usually kept in oil
- react readily with water to produce alkaline metal hydroxide and hydrogen gas
47
grp 1 reacting with grp 17
- react vigorously
- reaction rate increases down group 1
- as atoms get larger
- electron further away from nucleus
48
physical properties of grp 17
- density and metling point increases down the group
49
reactivity of grp 17
- decreases down the group
- electron affinity decreases
- atomic radius increases
-
50
halogen displacement reactions with halide ions
- occurs when a more reactive halogen displaces a less reactive halogen
- reactivity of grp 17 increases as you move up the group
- when halide ion is above the halogen in compound, it is replaced
- when Cl- replaces Br - solution becomes orange
-
51
Transition elements
-incomplete d-subshells or that can form atleast one stable cation with incomplete d-subshells
52
what d-block elements arent transition metals
Sc, Zn
53
what do transition elements have that other metals don't
- variable oxidation states
54
why transition elements form complex ions
- due to their variable oxidation states
55
what is a complex ion
molecule or ion, consisting of a central metal atom or ion, with anumber. ofmolecules or ions surrounding it
56
what is a ligand
- molecule or ion surrounding the central metal atom or ion
- due to different oxidation states of central metal ion, a different number and wide variety of ligands can form bonds with the transition elements
57
transition elements are able to form ____ compounds
colored
58
why do transition elements act as catalysts?
- since they have variable oxidation states
- during catalysis, the transition element can change to various oxidation states by gaining electron from or donating electrons to reagents withing the reaction
59
magnetic properties of transition elements
- diamagnetic, paramagnetic or ferromagnetic
- transition metals exhibit properties depending on their electronic configurations
- result of unpaired electron in the trainsition metal atom or ion
60
common ligands
H2O
NH3
Cl-
CN-
OH-
61
co-ordinate number
- number of coordinate bonds to the metal ions
62
naming complexes
prefix for number of ligands / ligand name/ element / oxidation number
63
transition metals as catalysts
- heterogeneous catalyst as they can provide a surface for reaction
- use 3d and 4s electrons to form weak bonds to reactant molecules
64
why do transition metals exhibit colors?
- d-orbital have the same energy in an isolated atom, but split into two sub-levels in a complex ion.
- electrical field of ligands cause the d orbital in complex ions to split so that the energy of an electron transition between them corresponds to a photon of visible light
65
what does the color exhibited by complex ion depend on?
- color observed is complementary to the color absorbed
- color depends on the identity of the metal ion, oxidation state of metal, and identity of the ligand
- ions with higher charge and ligands with greater charge density produce a greater split in the d-orbital
- spechtochemical series arranges ligans according to the energy seperation between the two sets of d orbitals.
