Notes – Unit 2 Section B / Chapter 5 Earth’s Mineral Resources / Electron structure

Notes – Unit 2 Section B /

Chapter 5Earth’s Mineral

Resources / Electron structure

Earth’s composition

1) Atmosphere: Gaseous part of earth. Provides nitrogen, oxygen, neon, argon, water vapor, and carbon dioxide



2) Hydrosphere: Liquid portion composed of water and dissolved minerals and salts.




Earth’s composition1) Atmosphere: Gaseous part of

earth. Provides nitrogen, oxygen, neon, argon, water vapor, and carbon dioxide


3) Lithosphere: solid part of earth. Greatest variety of chemical resources. Includes petroleum and metal-bearing ores.





Ore – Naturally occurring rock or mineral that can be mined


3) Lithosphere: Includes petroleum and metal-bearing ores.

Ore – Naturally occurring rock or mineral that can be mined.

Minerals -




Minerals – Naturally occurring solid compounds containing elements or groups of elements




Minerals – Naturally occurring solid compounds containing elements or groups of elements

Granite – Rock (Ore)

Quartz - mineral

Metals, Nonmetals and reactivity

• Metals are not generally found in their pure state.

• They react and combine with other elements to form ionic and covalent compounds.

• Their reactivity depends on their chemical families and properties

Types of Atomic particles

• Atoms - The building blocks of matter.

• Particles –Electrons (-) e-

Protons (+) p+

Neutrons (no charge) no• Nucleus Protons and neutrons

Atomic Particles

• Sodium =Atomic Mass =

Mass

# = - #(rounded) mass =

p+ =e- =no =

• Chlorine =Atomic Mass =

Mass


p+ =e- =no =

Atomic Particles

• Sodium = Na Atomic Mass = 22.99

Mass 23

# = 11 - # 11(rounded) mass = 23

p+ = 11e- = 11no = 12

• Chlorine =Atomic Mass =

Mass


p+ =e- =no =

Atomic Particles

• Sodium = Na Atomic Mass = 22.99

Mass 23

# = 11 - # 11(rounded) mass = 23

p+ = 11+e- = 11-no = 120

• Chlorine = ClAtomic Mass = 35 .453

Mass 35

# = 17 - # 17(rounded) mass = 35

p+ = 17+e- = 17-no = 180

Electrons

Electrons – Negatively charged particles orbiting the nucleusPeriods – Show the number of electron shells or orbitals. Horizontal columns.Families – Vertical columns. Show the number of outside electrons.

Electrons

Electrons – Negatively charged particles orbiting the nucleusPeriods – Show the number of electron shells or orbitals. Horizontal columns.Families – Vertical columns. Show the number of outside electrons.

1

23 4 5 6 7

8

1

2

3

4

5

6

7

Valence electrons

Periods

Electrons

Valence Electrons – The electrons found in the outside energy levelThe number of valence electrons determines many properties of that elementThe number of valence electrons make up families

Valence electrons

Electrons

Valence Electrons – The electrons found in the outside energy levelThe number of valence electrons determines many properties of that elementThe number of valence electrons make up families

Valence electrons

Honors ChemChapter 5

Electrons in Atoms

Atomic models• Dalton – 1st Atom – no internal

structures

• Thomson – Plum pudding, electrons in sphere of positive

• Rutherford - small dense positive nucleus with electrons around nucleus

• Bohr – electrons in circular orbitals with fixed distances from nucleus. Lowest energy inside highest energy furthest from nucleus

Atomic models

• Schrodinger – mathematical equation for location. Electron cloud model

• Quantum Mechanical model – modern description from Schrodinger’s mathematical equations

Atomic models

• Schrodinger – mathematical equation for location. Electron cloud model

• Quantum Mechanical model – modern description from Schrodinger’s mathematical equations

Electron configuration

• Quantum – the amount of energy required to move an electron from its present energy level to the next higher one

• Energy levels – region in space around nucleus where an electron is likely moving

Electron configuration• Quantum – the amount of

energy required to move an electron from its present energy level to the next higher one


• Energy levels are labeled by principle quantum numbers (n).

• n = 1,2,3,4, etc


Electron configuration• Quantum – the amount of

energy required to move an electron from its present energy level to the next higher one



• n = 1,2,3,4, etc• The higher the energy level

occupied by an electron, the more energetic it is



• n = 1,2,3,4, etc• The quantum # equals the

number of sublevels for each principle energy level

• The formula 2n2 equals the maximum number of electrons allowed in a principle energy level

• Shell n 2n2

• K 1 2(1)2=• L 2 2(2)2=• M 3 2(3)2=• N 4 2(4)2=

Energy levels


• n = 1,2,3,4, etc• The quantum # equals the

number of sublevels for each principle energy level

• The formula 2n2 equals the maximum number of electrons allowed in a principle energy level

• Shell n 2n2

• K 1 2(1)2=• L 2 2(2)2=• M 3 2(3)2=• N 4 2(4)2=

Electrons

• Electrons are found in orbitals or shells

• The electrons with the highest energy are on the outside shell.

• The electrons with the lowest energy are on the inside

Electrons

• Each shell can only hold a certain amount of electrons

• The electrons with the highest energy are on the outside shell.

• The electrons with the lowest energy are on the inside

# of electrons in orbitals

ShellK

LMN

n 2n2

1234


ShellK

LMN

n 2n2

1 2234


ShellK

LMN

n 2n2

1 22 834


ShellK

LMN

n 2n2

1 22 83 184


ShellK

LMN

n 2n2

1 22 83 184 32


ShellK

LMN

n 2n2

1 22 83 18 (8)4 32 (8)

Electron shells

• Electrons fill from the inside out

• The first shell has 2, the second shell has 8, the third shell has 18(8)

• Octet rule – The maximum number of electrons in the outside shell is 8. The exception is shell # 1 with 2.

Electron shells

• The atomic number of an element tells the number of protons and electrons

• Electrons are drawn in shells in pairs from inside out.

• The outside are valence electrons

Valence electrons

1

23 4 5 6 7

8

Electron shells




N

Atomic # = 7 1st = 2Period = 2 2nd = 8

3rd = 8

Electron shells




S

Atomic # = 16 1st = 2Period = 3 2nd = 8

3rd = 8

1

23 4 5 6 7

8

Electron dot diagram

• Shows each element with its valence electrons only

• Outside electrons are placed in pairs

• Dots show outside electrons

S





• Elements will lose or gain e- to become stable and reach the octet (8 or 0)

Electron dot diagram• Shows each element with its valence

electrons only• Outside electrons are placed in pairs• Dots show outside electrons

• Elements will lose or gain e- to become stable and reach the octet (8 or 0)

Elements that lose e- become positive , those that gain e-

become negative

Li =

C =

O =

1

23 4 5 6 7

8


electrons only• Outside electrons are placed in pairs• Dots show outside electrons• Elements will lose or gain e- to become

stable and reach the octet (8 or 0)

Elements that lose e- become positive charged , those that

gain e- become negative charged

• Dots are placed R-L-T-B and then repeat pattern in pairs

Al =

N =

Cl =

Ar =


electrons only• Outside electrons are placed in pairs• Dots show outside electrons• Elements will lose or gain e- to become

stable and reach the octet (8 or 0)

Elements that lose e- become positive charged , those that

gain e- become negative charged

• Dots are placed R-L-T-B and then repeat pattern in pairs

Al = lose/gain _____ charge = ___

N = lose/gain _____ charge = ___

Cl = lose/gain _____ charge = ___

Ar = lose/gain _____ charge = ___





S

Electrons and Periodic Table

• Valence electrons in atoms show reactivity

• Atoms will lose or gain electrons to reach the octet (8) rule level.

• Inert (Noble) Gases have 8 valence electrons (2 for He) and are stable



• Nonmetals –Groups 5a to 7a will gain electrons to become stable

• Metals – Groups 1a to 3a will lose electrons to become stable



• Nonmetals –Groups 5a to 7a will gain electrons to become stable

• Metals – Groups 1a to 3a will lose electrons to become stable

Orbitals

• Atomic orbitals – regions around the nucleus within which the electrons have the highest probability of being found

• Orbitals are sublevels which correspond to different shapes.

• s are spherical, p are dumbell shaped, etc

• Sublevel orbitals can contain differing numbers of orbitals.

• s = 1 orbital• p = 3 orbitals• d = 5 orbitals• f = 7 orbitals

Electron Arrangements in Atoms

Electron configuration – The ways in which electrons are arranged around the nuclei of atoms3 Rules:•1) Aufbau principle – electrons enter orbitals of the lowest energy first•2) Pauli exclusion principle

Electron Arrangements in Atoms

Electron configuration – The ways in which electrons are arranged around the nuclei of atoms3 Rules:•1) Aufbau principle – electrons enter orbitals of the lowest energy first•2) Pauli exclusion principle

Electron Arrangements in Atoms3 Rules:•2) Pauli exclusion principle – atomic orbitals can hold no more than two electrons•3) Hund’s rule – When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins.

Calculating Electron Configurations

• The shorthand method for writing electron configurations of an atom involves writing the energy level and symbol for every sublevel occupied by an electron and indicating the number of electrons in each sublevel with a subscript


• The shorthand method for writing electron configurations of an atom involves writing the energy level and symbol for every sublevel occupied by an electron and indicating the number of electrons in each sublevel with a subscript.

Calculating Electron Configurations• The shorthand method for

writing electron configurations of an atom involves writing the energy level and symbol for every sublevel occupied by an electron and indicating the number of electrons in each sublevel with a subscript.

• Nitrogen – Atomic # 7• Shorthand = 1s22s22p3

• Write the electron configuration for Phosphorus,

• atomic # 15


Write the electron configuration for Phosphorus, atomic # 1515 e-


Write the electron configuration for Phosphorus, atomic # 1515 e-

Calculating Electron ConfigurationsConfiguration = 1s22s22p63s23p3 = 15 electrons

• Write the electron configuration for Carbon, Argon, and Nickel


• Write the electron configuration for Carbon (6), Argon (18), and Nickel(28)


Write the electron configuration for Carbon, atomic # 66e-


Write the electron configuration for Argon, atomic # 1818e-


Write the electron configuration for Nickel, atomic # 2828e-


Physics and the Quantum Mechanical Model

• Quantum mechanical model grew out of the study of Light

• Light consists of waves and particles

• Wave principles:• Amplitude = wave

height• Wavelength ( ) =

distance between crests


• Quantum mechanical model grew out of the study of Light

• Light consists of waves and particles

• Wave principles:• Amplitude = wave

height• Wavelength ( ) =

distance between crests


• Wave principles:• Amplitude = wave height• Wavelength ( ) = distance

between crests• Frequency ( ) = wave

cycles past a given point per unit time– Cycles/second = hertz (Hz) s-1

• Speed of light constant


• Wave principles:• Amplitude = wave height• Wavelength ( ) = distance

between crests

• Frequency ( ) = wave cycles past a given point per unit time– Cycles/second =

hertz (Hz) s-1



• Frequency ( ) = wave cycles past a given point per unit time– Cycles/second = hertz

(Hz) s-1




C = 2.988 x 108 meters/second• (186,000 miles/second)



C = 2.988 x 108 meters/second• (186,000 miles/second)• Wavelength and frequency are

inversely proportional

– High = Low– Low = High

• All electromagnetic waves travel at same speed = 3.0 x 108 m/s


• Wavelength and frequency are inversely proportional



• Light = Quanta of energy calledPhotons

• Atomic Spectra – When atoms absorb energy, electrons move into higher energy levels. These electrons than lose energy by emitting light.

High

High Low

Low

























Chemical Bonds – Chemical forces which hold atoms together and form complete

electrons shellsIon – is an atom or group of

atoms that have an electric charge.Atoms become charged when they lose or gain electrons to become stableAtoms which lose electrons become positive (+) Atoms which gain electrons become negative (-)

Chemical Bonds – Chemical forces which hold atoms together and form complete

electrons shells• Ionic Bond – Positive (+) and

negative (-) ions attract – formed by the transfer of valence electrons.Na 11p+ 11p+

11e- 10e- = +1 Na

Cl 17p+ 17p+17e- 18e- = -1 Cl

Na +1 + Cl-1 = NaCl o

Documents

Notes – Unit 2 Section B / Chapter 5 Earth’s Mineral Resources / Electron structure