Chem Test #4 Flashcards
Electronegativity _____ left to right
Increases
Z eff (effective nuclear charge) ___________ top to bottom
Stays the same
Electron affinity becomes ___________ top to bottom
Less negative
Electronegativity __________ top to bottom
Decreases
Inner core electrons repelling
The valence electrons do not experience the full attraction of the nuclear charge because the inner core electrons reduce the positive nuclear charge (shielding effect)
The electrons in different levels are closer to each other than the nucleus so they repel
Net nuclear charge formula
Z eff = Z - S
Z is the atomic number
S is the number of inner core electrons (non valence)
Parent atom vs ion
The cation ion will be smaller than the parent ion because it loses electrons and the valence electrons are more strongly attracted to the nucleus
The anion is the opposite
Size of cations across a period
When the cations all have the same electron configuration the nuclear charge increases across a row. The increased attraction between the electrons pulls the outer shell closer. So… the atomic radius decreases.
Ionization Energy
The minimum amount of energy required to remove an electron from an atom
Ionization Energy
The minimum amount of energy required to remove an electron from an atom
IE increases from left to right because….
The valence electrons get closer as the Z eff increases so there is a greater force of attraction.
IE decreases from top to bottom because….
The Z eff stays the same down a group so the nuclear charge doesn’t affect it.
The atomic radii is larger and makes them easier to remove.
Exceptions to the IE rule
Boron and oxygen
Boron: The outer most electron in Boron is further away from the nucleus in the 2p orbital
Oxygen: In an orbital with a second electron and experiences repulsion
Electron affinity
The energy change that occurs when an electron is acquired by a neutral gaseous ion
It releases energy and is measured in kJ mol-1
Noble gases do not have ea
Which group has the most negative ea
Group 17 because they love to attract electrons
What groups has the max ea value and why
Group 2 because the electron is being added to a new orbital
Group 15 because the electron is being added to a orbital with one electron already with repulsion
Electronegativity
The relative attraction that an atom has for a shared pair of electrons in a covalent bond
Why the electronegativity increase from left to right and decrease up and down
Increase because the Z eff increases and the atom has an increased attraction
Decrease because the radii increases and the bonding electrons are further from the nucleus
Physical properties of alkali metals
Good conductors
Low densities
Low melting points
Soft
Melting and boiling points decrease down a group
Chemical properties of alkali metals
Ionic compounds
Solids
Reactions with water become more vigorous down a group
Reactions with water and alkali metals
They all produce metal hydroxide and hydrogen gas
Basica solution produced
H20 is the oxidizing agent, metal is reducing one
Physical properties of group 17
Diatomic molecules
F and Cl are gases, Br is a liquid, I and At are solids
Melting/Boiling points increase down the group
Colours of group 17
F -pale yellow
Cl - yellow green
Br - red brown
I - purple
Chemical properties group 17
Very reactive
Reactivity decreases down a group
Reaction of halogens with group 1 metals
Alkali metals lose one electron (oxidized)
Halogens gain one electron (reduced)
Form an ionic bond
Halogens bonding with ionic bond with a metal and a halogen
Down the group the oxidizing ability decreases
The smaller (radii) halogen displaces the larger halogen
Takes on the colour of the one that was displaced
Determining between I2 and Br2
Add hydrocarbon solvent to both solutions, I2 will become purple
Period 3 oxides
Ionic character of period three oxides decrease from left to right
Become more ionic down a group
Oxides only conduct electricity in the molten state
Period 3 oxides basic and acidic
Basic - Na2O, MgO
Amphoteric - Al2O3
Acidic - SiO2, P4O10 / P4O6 , SO3 / SO2, Cl2O7 / Cl2O
Oxidation state of an element
0
Oxidation states in an ionic compound
Equal to their charges
Definition of transition metals
An element that forms at least one stable cation with an incomplete 3d sublevel
Why there is a small decrease in atomic radii
The nuclear charge increases because 1 proton is added but this is offset by the addition of the inner 3d electron
Physical properties of transition metals
-high electrical and thermal conductivity
-high melting point
-malleable
-high tensile strength
-ductile
Delocalized electrons
The electrons involved in metallic bonding that are free to move easily from one atom to the next throughout the metal and are not attached to a particular atom
High conductivity
Magnetic properties of transition metals
Since they unpairs of spinning electrons, they have magnetic properties
Scandium forms a ______ solution
Colourless
What is a catalyst
Substance that speeds up the rate of a chemical reaction
Why transition metals have multiple oxidation states
Why transition metals have multiple oxidation states
Why does the 4d orbital split
It is because the ligands with the lone pair of electrons are in between the axises and experience less repulsion from the nucleus, and therefore have more energy
Why does these ligand compounds emit colour
An electron wants to go from the lower energy level to the higher one, this emits colour
Energy gap vs wavelength of light
the bigger the gap the smaller the wavelength, more energy in the wavelength, higher frequency in the wavelength
Colour absorbed vs emitted
Complemntary
How charge effects the colour
As the metal ion charge increases, the gap energy increases
The energy of the light absorbed increases
How indignity of the metal effects the colour
More d electrons means higher gap energy
How type of ligand effects the colour
We replace some of the other ligand with a new one and see what colour is produced
If the colour has a higher energy, the new ligand has a greater splitting ability
