Unit 2 IB CHEM

AMU & Location of Electron, Neutron, and Protons

 

Isotopes can be written in two ways:

Isotopes

• Isotopes = atoms of the same element with a different number of neutrons and mass

Properties of Isotopes

• Chemical Properties of isotopes of the same element are the same 

  • Because chemical behavior is associated with electrons not neutrons

• Physical Properties of isotopes of the same element are different

  • Physical properties are based on mass

  • Densities, MP, BP, etc…will be different

  • Some isotopes are radioisotopes (radioactive)

Mass Number

• The mass number is the number of protons and neutrons in the nucleus.

• Mass Number is NOT the actual mass of an atom or element

Average Atomic Mass

• Average atomic mass is an AVERAGE of ALL of that element in the universe.

  • Atomic Mass is expressed in Atomic Mass Units → amu

• The periodic table lists average atomic mass based on mass and percent (%) abundance of each naturally occurring isotope of an element.

Average atomic mass is a weightedaverage

Abundance

• Abundance refers to the amount of each isotope

  • The average atomic mass of an element is going to be closest to the most abundant isotope

Average Atomic Mass Formula

(Mass x Abundance % ) + (Mass x Abundance % ) + …

               100                   100

Valence electrons

• The valence electrons are the outermost electrons (the highest energy level).

• Valence electrons determine chemical properties.

•  Every element wants to have a full octet = 8 Valence e^-

  • Exceptions: 

    • H: Full with 2 valence e^-

    • B: Full with 6 valence e^-

  • Elements will gain or lose valence electrons to get a full octet, this is how ions are formed

Periods

• There are 7 periods

• Elements in the same period have the same number of electron shells/energy levels

  • The Horizontal rows

Groups/Families

• The vertical columns of the periodic table are called GROUPS or FAMILIES.

• The elements in any group of the periodic table have similar physical and chemical properties.

Identifying Electrons in an Atom

• 4 Unique Quantum Numbers are used to identify each electron in an atom

  • Principal Quantum Number- Principal energy level (n)

  • Second Quantum Number- Sublevel (l)

  • Magnetic Quantum Number- Orbital (m_l)

  • Spin Quantum Number- Spin (m_s)

Principal Quantum Number

• Principal Energy Level (Energy Level)

  • positive integer

  • n = 1, 2, 3, 4, ...

• larger numbers are farther away from the nucleus - higher energy level

  • Energy level 1 = max 2 electrons

  • Energy level 2 = max 8 electrons

  • Energy level 3 = max 18 electrons

  • Energy level 4 = max 32 electrons

Second Quantum Number

• Sublevel - type or Shape of Atomic Orbitals

• You can remember this with the saying:

  • Some People Do Fine

 

Principal Energy Level	Sub-levels
1	s
2	s, p
3	s, p, d
4	s, p, d, f

Shape of S Orbital (Must know how to draw)

Shape of P orbital (Must know how to draw)

Shape of d orbital and f orbital (not required to draw)

Magnetic Quantum Number

• Represents the number of orbitals in the sublevel (which orbital are they in?)

• Each orbital can hold up to 2 electrons

 

Sub-level	Number of orbitals	Number of electrons
s	1	2
p	3	6
d	5	10
f	7	14

 

Principal Energy Level	Sub-level	# of orbitals in sub-level	# of electrons in  Energy Level
1	s	1	2
2	s	1	2
	p	3	6
3	s	1
	p	3
	d	5
4	s	1
	p	3
	d	5
	f	7

Not a flashcard

 

Principal Energy Level	Sub-level	# of orbitals in sub-level	# of electrons in  Energy Level
1	s	1	2
2	s	1	2
	p	3	6
3	s	1
	p	3
	d	5
4	s	1
	p	3
	d	5
	f	7

Not a flashcard

Spin Quantum Number

• Orientation of Electron (Spin)

• Each orbital can only hold 2 electrons

• Electron spin in the same orbital is opposite

• ms = + ½ or - ½

Rule #1: Aufbau Principle

Start filling electrons at the lowest energy level. One electron at a time

Rule #2: Pauli Exclusion Principle

A maximum of 2 electrons can be in one orbital.  They must have opposite spins.

Rules #3: Hund’s Rule

Electrons fill each orbital in a sublevel before they start pairing up

Noble Gas Configuration

• Find the noble gas that is one row above your element.

• Write that noble gases symbol in [  ].

• Then write the remaining sublevels that are filled by your element.

Valence Electrons

Valence electrons are electrons held in highest energy level – typically the last s or s and p sub-levels

Ground State and Excited State

• When the electrons in an atom become excited by absorbing energy from their surroundings, they jump to new higher energy levels

• Electrons in the lowest energy level possible are in the ground state.

• Electrons in ANY higher level are in an excited state.

Relation between length of a wave and its energy

• The length of the wave is inversely related to its energy

Not a flashcard

Not a flashcard

Return to the Ground State

• When electrons fall back, a specific amount of energy is emitted based on the distance between electron orbitals.

• Orbitals further from the nucleus are closer together.

Return to the Ground State - 1^st energy level

• When electrons return to the 1^st level, the lines appear in the ultraviolet region (Lyman series) since it is the largest energy change

Return to the Ground State - 2^nd energy level

• When electrons return to the 2^nd level the lines appear in the visible region (Balmer series). 

Return to the Ground State - 3^rd energy level

• When electrons return to the 3^rd level, the first series of lines in the infrared region (Paschen series) is produced.

Data that line spectra’s give

• 1.  Electrons exist in distinct energy levels because each line represents the energy emitted as an electron falls from excited to ground state

2.  Energy differences are smaller between levels that are further from the nucleus because lines are closer together on the higher energy end of the spectrum

Alkali Metals

• Properties

  • Metallic properties

  • Soft

    • Can be cut with knife

• Reactivity increase down the group

• highly reactive with O₂, H₂O

• produce alkali (base solns)

• Electron Configuration Xs¹

• Lose 1e^- → Noble Gas

Alkaline Earth Metals

• Properties

  • harder, denser, stronger, and higher melting point than alkali metals

• react with O²

• Reactivity increases down

• Electron Configuration Xs²

• Lose 2e^- → Noble Gas

Transition Metals

• Properties

  • metals, harder, denser, higher melting point (except Hg)

  • not as reactive as groups 1&2

• Inner transition elements

  • Lanthanides 58-71

  • Actinides 90-103

Halogens

• Combine with metals to form salts

  • NaCl

• most reactive nonmetals

• Reactivity decreases down

• Electron Configuration

• Xs² Xp⁵ (gain 1e- → Noble Gas)

Noble Gases

• Inert Gases

  • Unreacted

  • Already Full Shell ( 8 Valence e^-)

  • Stable

• Electron Configuration Xs² Xp⁶ 

  • except He - 1s²

• Other elements want the same configuration

Nuclear charge (trend)

• Across – increases

• Down – increases

Nuclear charge – the charge in the nucleus or the number of protons

Valence Electrons (trend)

• Valence Electrons – electrons in the outermost energy level (s and p orbitals only!)

  • Across – increase

  • Down – stays the same

Average Atomic Mass (trend)

• Average Atomic Mass – weighted average based on mass and percent abundance of each naturally occurring isotope (Remember there are exceptions to this (Te/I ) but it is generally true)

  • Across – increases

  • Down - increases

Shielding Effect

The shielding effect is the reduction of attractive force between the nucleus and its valence electrons due to the blocking effect of the shielding electrons

• Across the period, stays the same

  • Electrons are added to valence shell 

  • The number of “inner” electrons”remain the same

  • Same number of electron shells

• Down the group, it increases

  • Energy level are being added

  • Inner electrons increase

• Atomic size increases

  • More distance between valence shell and the nucleus

  • Bigger distance → less attraction = higher shielding effect

Atomic Radius

• Half the distance from center to center of like atoms

• Across the period, decreases

  • The number of protons increases

• Down the group, increases

  • Energy levels are being added, leads to bigger atomic size

  • Shielding increases

Ionization Energy

The amount of energy needed to remove an electron from an atom

• Across the period, increases

  • Greater the attraction between the valence electrons and the nucleus

• Down the group, decreases

  • Shielding increases, thus atoms get bigger

Electronegativity

• The ability for an atom to attract electrons to itself when it is bonded to another atom

• Across the period, increases

  • More protons are being added → Greater attraction for electrons

• Down the group, decreases

  • Increased in shielding, which blocks nuclear attraction

Electronegativity Exceptions

• Noble gases have an electronegativity of zero because they are inert or unreactive.

• Already have a full valence shell

  • Octet rule - 8 valence electrons

Reactivity

• Reactivity is how readily an element will react with another element. 

• Metals and Non-metals don’t follow the same trends.

Reactivity in Metals

• Across the period, decreases

  • Nuclear charge increases 

• Energy levels stay constant, so ionization energy increases.

• Down a group, increases

  • Nuclear charge increases, 

  • Energy levels (shielding effect) increases, so ionization energy decreases. 

Reactivity in Nonmetals

• Non-metals react by gaining electrons. 

They have a high electronegativity , so it's fairly easy for them to gain electrons

• Across the period, increases

  • Energy levels stay constant, so electronegativity increases. 

• Down the group, decreases

  • Energy levels (shielding effect) increases, so electronegativity decreases.

Isoelectronic Series

• A series of ions/atoms with the same electron configuration (valence isoelectronic)

• O^2-, F^-, Ne, Na⁺, Mg²⁺

• 1s²2s²2p⁶

Ionic Radius - Anions

• Nonmetals get larger when they become anions.  

• This is because they are gaining electrons, so their electron cloud gets bigger while the number of protons stays the same

• With an additional electron, there is also additional electron repulsion 

Ionic Radius - Cation

• Metals get smaller when they become cations.

• This because they are losing electrons, so their electron cloud get smaller

• The nucleus exerts a stronger pull on fewer negative charges typically an energy level is lost.