THE PERIODIC TABLE & QUANTUM THEORY

Unit 3

QUANTUM THEORYLight and Electrons

Light; is it a wave or particle?

YES!Evidence for the Dual Nature of Light

Wave Properties of EMS

■ Velocity (v): wave speed, all forms of electromagnetic radiation travel at the speed of light (c = 3*108

m/s)■ Frequency(ν): number of peaks or

cycles in a second (1/s, s-1, Hz)■ Wavelength (λ): distance between

two equivalent points on consecutive waves (m, nm, cm, …)

■ Amplitude(ψ): height from peak to baseline (m, nm, cm, …)

Types of Electromagnetic Radiation

EM Spectrums

■ All types of radiation travel at the speed of light (c) = 3.0 x 108 m/s

■ Waves with a long wavelength have a low frequency and low energy

■ Waves with short wavelengths have high frequencies and high energy

■ Direct Relationship -> Energy & Frequency■ Indirect (inverse) Relationship -> Wavelength & Energy,

Wavelength & Frequency

High EnergyLow Wavelength

Low EnergyHigh Wavelength

Frequency (s-1) High FrequencyLow Frequency

Use the EM Spectrum to answer the following questions.■ 1) Which type of wave has the greatest frequency?

■ 2) Compare the colors of the visible spectrum, which has the lowest energy?

■ 3) Compare x-rays to microwaves, which one has the highest frequency?

Compare the energy of the following waves.

A

B

C

Definitions

■ Ground state: all electrons are in their lowest possible energy levels (stable)

■ Excited state: electrons have absorbed energy and moved to a higher energy level

■ Quantum: the exact amount of energy needed to move an electron from one energy level to another

■ Photon: a packet of electromagnetic radiation carrying a quantum of energy

What is the photoelectric effect?

DEF: When light of a certain frequency shines on a metal, as a result an electron is emitted from the metal

Warm Up – 9/20

■ What happens when electrons fall back down to the ground state?

■ Using the Bohr Model (On your reference table) What type of energy is released when an electron transitions from n=3 to n=1?

Why is Bohr’s model of the atom inadequate?■ It only works for hydrogen!

Assumptions of the Bohr Model

■ Electrons circle the nucleus in fixed energy levels (orbitals)■ Electrons transition between energy levels when they gain or

loose energy■ Energy level increase in energy as they move farther from

the nucleus

Atomic Emission Spectrum

■ Atomic Emission Spectrum = colors of light emitted by excited electrons as they move to a lower energy level; unique to every element

■ Ex. Neon emits light when excited.■ DEMO TIME

Atomic Emission Spectrum Video

More Definitions

■ Absorption spectrum: colors of light absorbed by electrons as they move to a higher energy level, dark bands

■ Spectroscopy: method of studying the interaction between matter and electromagnetic radiation

How do we know what the Sun is made of?

Let’s Try It!

Using the Bohr Model on the back of the reference table answer the following:

■ 1) When the electron goes from n=4 to n=1 what portion of the EM spectrum is viewed?

■ 2) When the electron moves from n=6 to n=3 what portion of the EM spectrum is viewed?

■ 3) What color is emitted when the electron moves from n=3 to n=2?

Quantum Theory

■ Proposed by Max Planck (1900)■ What is Quantum theory?

– (particle nature of light) Energy is not emitted continuously but in small packets known as quantaPlanck’s constant = 6.626*10-34J*s

Using the PT to determine Electron Configurations & Orbital Notations■ Orbitals – 3-D region where electrons are

located■ Sublevels – energy level in which the

orbitals are located■ The periodic table can be split into

“sections” called sublevels. These sublevels are the “s”, “p”, “d” and “f” sublevels

■ You will need to memorize the location of sublevels

Sublevels

■ s sublevel = 2 e-, 1 orbital■ p sublevel = 6 e-, 3 orbitals■ d sublevel = 10 e-, 5 orbitals■ f sublevel = 14 e-, 7 orbitals

■ Each orbital can only have 2 e- max!

Electron Configuration

■ Format = #letter#

– # = energy level– Letter = sublevel– Superscript # = # of electrons

■ Use the periodic table as your guide

■ Important Rules with the “d” and “f” block– d starts at 3 (1 minus the row/period #)– f starts at 4 (2 minus the row/period #)

Electron Configuration

■ Example for Fe (Iron)

Practice Problems

1. O

2. K

3. Ni

4. Te

Noble Gas Electron Configuration

■ This is a short-hand version of electron configuration

■ Format: [X]…….– X = noble gas that comes directly before the element

numerically– …. = the rest of the electron configuration from that

noble gas to the element

Practice Problems

1. O

2. K

3. Ni

4. Te

Warm Up – 2/12

■ Write the electron configuration for magnesium.

■ What element is indicated by the following electron configuration?

1s22s22p63s23p64s23d7

Electron Configuration and Orbital Diagrams■ Rules for determining electron

locations– Electrons must fill the lowest

energy level first– There must be an electron

within each orbital before they pair up

– Electrons that occupy the same orbital must have opposite spin directions

Orbital Diagrams

Practice Problems

1. O

2. K

3. Ni

4. Te

Warm Up – 2/13

■ Last Minute Study for your quiz– Light (compare energy, frequency, and Wavelength)– Bohr Model (tracing the lines)– Electron Configurations– Noble Gas Configurations– Orbital Notation

THE PERIODIC TABLEA chemist’s best friend!

HISTORYHow did it end up looking

like the one in our classroom?

Early History

■ Newland’s Law of Octaves (1863)– When arranged by atomic mass, properties of the elements

repeated every eighth element

■ Mendeleev’s Periodic Table (1869)– Arranged elements by atomic mass and elements in the same

column had similar properties

Mendeleev

■ What was significant about Mendeleev’s table?– He left empty spaces for “undiscovered” elements. He used his

table to predict the properties of these elements (Scandium (Sc), Gallium (Ga), Germanium (Ge)) and he was right.

■ What questions were raised by the arrangement of the periodic table?– Why did some of the elements not follow the trend?– Why are the elements periodic?– Genius of Mendeleev

Modern Periodic Table

■ Moseley’s discovery of atomic number (1913)– Lead to a rearrangement of the period table, by

increasing atomic number; columns of elements had similar properties

■ Modern Periodic Law– Properties of the elements are a periodic function of

their atomic numbers

ARRANGEMENT OF THE MODERN PERIODIC TABLE

Metals

■ Located to the left of the staircase + Aluminum

■ General Properties: – Good conductors of heat and electricity– Few valence electrons– Malleable - can be hammered or beaten into sheets – Ductile - can be drawn, pulled, or extruded through a small

opening to produce a wire– Shiny– High melting & boiling points– Mostly solids (exception: Hg, Mercury)

Nonmetals

■ Located to the right of the staircase

■ General Properties: – Poor conductors of heat and electricity– Mainly gases (exception Bromine)– Dull brittle solids– Low melting & boiling points– 5-8 valence electrons

Metalloids

■ Located along the staircase except Aluminum & Astatine

■ General properties: – Solids– Properties of metals and

nonmetals– Semiconductors

Periodic Table Organization

■ The Periodic Table is arranged by groups and periods

■ Groups = Families– You will need to know the name of each family on the periodic

table.

■ Groups = vertical columns

■ *Key Idea - Elements in the same group have the same number of valence electrons, which helps determine reactivity.

Periodic Table Organization

■ Each horizontal row on the periodic table = Period

■ The period helps to determine how many energy levels each element has.

■ Example: Sodium is in the 3rd horizontal row. We call this the 3rd

period, which means it has 3 energy levels.

Periodic Table Families

Main Group/Representative Elements

■ Alkali Metals: Group 1, most reactive metals, not found free in nature, silvery, soft, react with water, one valence electron (Hydrogen is a NON METAL!!!)

■ Alkaline Earth Metals: Group 2, reactive (not found free), harder, denser, stronger than group 1, two valence electrons

Main Group Elements

■ Halogens: – Group 17– Most reactive nonmetals– “salt formers”– Gases (F, Cl)– Liquid (Br)– Solids (I, At)– 7 valence electrons

More Main Group Elements

■ Noble Gases: Group 18, stable, quite unreactive, although compounds have been formed with some, eight valence electrons (except Helium which has 2 valence electrons)

Warm Up – 2/14

■ To what family do the following belong. Is it a metal, nonmetal, or metalloid?– Ca– I– U– Na– He– Ag– Dy– F

d – Block Elements

■ Transition Metals:– Located in the d block– Metallic properties– Less reactive, harder, stronger than Groups 1 & 2

Inner Transition/Rare Earth Metals

■ Lanthanides: discovered in early 1900’s, atomic numbers 58-71, shiny metals, reactivity similar to group 2

■ Actinides: atomic numbers 90-103, all radioactive

Valence Electrons - (outer energy level, s and p sublevels)

Group 1 2 3-12 13 14 15 16 17 18

Number of Valence

Electrons1 2 2 3 4 5 6 7 8

Oxidation Number(Charge)

+1 +2 varies all +

+3 -3 -2 -1 04+−

+ lose electronsCations

- gains electronsAnions

Maximum amount of valence electrons = 8

Cations and Anions

■ Cations have a POSITIVE charge because they LOSE electrons (these are usually metals)

■ Anions have a NEGATIVE charge because they GAIN electrons (these are usually non metals)

Valence Dot Diagrams

Valence Dot Diagrams

Practice Problems: Calcium:Nitrogen:Helium:Xenon: Selenium

PERIODIC PROPERTIES AND TRENDS

Different Trends

■ Atomic radius■ Ionic radius■ Reactivity■ Ionization energy■ Electronegativity

Atomic Radius

■ Definition: half the distance between nuclei of identical atoms bonded together

■ Trend Across a row: decreases due to increasing nuclear charge (electron cloud pulled in tighter to nucleus)

■ Trend Down a group: increases due to an increase in the number of energy levels

Atomic Radius Continued

Which element is the largest element on the periodic table?

WHY?

Apply It!

1) Bigger: Li or K?

2) Larger: C or F?

Ionic Radius

■ Definition:– (+) ion = cation: lost an electron, decrease in atomic radius Na >

Na+

– (-) ion = anion: gain an electron, increase in atomic radius Cl- > Cl

■ Trend: same as atomic radius for the same reasons– Across a row: decreases (increase in nuclear charge), the

increases as the first negative ion forms, then decrease again (increase in nuclear charge)

– Down a group: increases (increase number of energy level)

Apply It!

1) Larger: Ca or Ca+2?

2) Larger: Br or Br-1?

Reactivity■ Metals: most active is Francium

(largest, easiest to lose electrons)

■ Nonmetals: most active is Fluorine (smallest, hardest to lose electrons)

■ Metallic Behavior/Reactivity: increases left and down– Towards Fr

■ Nonmetallic Behavior/Reactivity: increases up and right– Towards F

■ Reactivity

Apply It!

1) More reactive: K or Cs?2) Less reactive: Cu or Ca?3) More reactive: Cl or I?4) Less reactive: C or O?

Ionization Energy

■ Definition: energy needed to remove the outermost electron from a neutral atom, metals generally have a low I.E. and nonmetals have a high I.E. (Also known as F.I.E.)

■ Trend Across a row: increases (increase in nuclear charge)

■ Trend Down a group: decreases (increase energy level, shielding)

Apply It!

1) Higher: Cl or I?

2) Higher: Na or S?

Electronegativity

■ Definition: the ability for an atom to attract electrons

■ Trend: Increases as you go across the periodic table. Decreases as you go down a family. (Omit the noble gases for this trend)

■ Fluorine is the most electronegative element (WHY do you think this is?)

Apply It!

1) Higher: C or N?

2) Higher: Cu or Fe?

Summarizing Trends

Warm Up – 9/19

■ Last Minute Study– Group/Family Names– Valence Electrons– Trends (electronegativity, ionization energy, reactivity,

atomic radius, ionic radius)

Warm Up – 2/19

■ Your test is Thursday 9/26■ Write the electron configuration and the noble gas

configuration for magnesium■ Rank the following in order of increasing ionization energy

– Te, S, Rb, Y

Warm Up – 2/21

■ Last Minute Study for Test!■ Study Guides Due