Honors Ch3 and Ch4 - Manasquan Public Schools...Ch. 3.1 The Atom is Defined 400 B.C. the Greek...

HonorsCh3 and Ch4

Atomic History

and the Atom

Ch. 3.1The Atom is Defined

400 B.C. the Greek philosopher

Democritus said that the world was made

of two things:

– Empty space and tiny particles called

atoms

Ch3.1 What are atoms?

Atoms are the smallest

part of an element that

still has the element’s

properties.

By 1700’s chemistry was defined by 3 Laws:

Law of the Conservation of Mass

Law of the definite Proportions

Law of Multiple Proportions

Law of Conservation of Mass

Law of Definite Proportions/Composition

Substances contain atoms in the same

ratio of mass.

Law of Multiple Proportions

Early 1800’s John Dalton

Came up with first atomic

theory that is the basis for

today’s theory.

Early 1800’s John Dalton’s Theory-proven

1. Every element is made of tiny, unique particles called atoms– These atoms cannot be destroyed, but

instead rearrange during a chemical change

2. Atoms of different elements can join to form molecules in constant whole number ratios.

John Dalton’s Theory-disproved

Atoms cannot be broken down into smaller particles.

Atoms of the same element are exactly alike in mass

Ch3.21897 JJ Thomson

Used a cathoray tube to

examine if atoms were made

of charged particles.

Opposite charges attract

1897 JJ Thomson

Discovered atoms are

made up of particles

with negative charges,

but little mass.

Called them electrons.

Electrons

Positive charges, not

known as protons yet

Thomson Model

1911-Rutherford

Put together a team of

physicists to performed the

gold foil experiment

1 out of 8000 alpha

particles were repelled.

1913-Rutherford

The experiment led to the discovery that atoms are mostly empty space. (expected)

It also discovered that atoms contain a positive dense nucleus which contained most of the mass of the atom.

Rutherford Model

Inside the Nucleus Particles were discovered.

Positive protons (1836 times more massive

than electrons)

– These practices identified the type of atom.

Neutral neutrons (1837 times more massive

than electrons)

– Kept the protons from repelling by producing strong

nuclear forces

Ch 3.3 Parts of an atom

Nucleus

–Proton

–Neutron

Cloud

–Electron

Subatomic Particles

Charged Particles

Nucleus

center of an atom

positively charged

makes up 99.9% of the atom’s mass

Protons

Charge (+)

Mass is equal to

1 atomic mass unit (u)

–1/12 mass of Carbon atom

Neutrons

Charge (0 net) - neutral

Mass is equal to 1amu

Determine stability of

nucleus

Electrons

Charge is negative (-)

Mass is equal to 0.001 amu

Atomic Number

Identifies # of protons

Determines the type of atom because

no two elements can have same # of

protons.

Mass of a single atom

Mass Number

# of

protons

# of

neutrons

Mass

#

Isotopes

Any atoms having the same number of protons but different number of neutrons.

thus they have different mass numbers.

Springfield Isotopes

Isotopes

Average Atomic Mass

Atomic mass is the mass of all

isotopes of a particular element

averaged together

Calculating Average Mass is based

on the abundance of each isotope

Atoms vs Ions

All atoms have the same

number of protons and

electrons.

They are neutral.

Charges cancel each other out.

Atom vs Ions

Ions are charged particles.

Form when atoms lose or gain electrons.

Form in order to have a full outer shell

Two Types.

Cations

Positively charged ions.

Form when atoms lose

electrons.

Form from metal atoms

Cations

# of protons greater

than # of electrons

More (+) than (-)

Na AtomNa+ Cation

Anions

Negatively charged ions.

Form when atoms gain

electrons.

Form from nonmetal atoms

Anions

# of protons less

than # of electrons

More (-) than (+)

Cl atom

Cl- Anion

Ch4.1 The Duality of Light Led to a New View of the Atom

Light has characteristics of both waves and

particles

All forms of radiation travel at the same

maximum speed of 3.00 x108 m/s

Wave description

Different forms of light are defined by their

unique wavelengths (ƛ) and frequencies (v)

Speed of light (c)= ƛ v

Electromagnetic Spectrum

Electromagnetic spectrum

includes light at all possible

frequencies and wavelengths

↑ wavelength, ↓ frequency

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Gamma Rays

highest energy and frequency.

Nuclear radiation

X-Rays

high ionizing energy radiation

Used for imaging

Ultraviolet Light (UV)

Ionizing energetic radiation

Can cause skin cancer

when over exposed

Sun Protection (SPF)

7 or less→no protection

8→Extra protection but still permits tanning

15→Offers total protection from burning

30→Totally blocks UV

ROYGBIV

Visible spectrum

We see red and orange the best

Blue and violet are the hottest colors, and

emits the most energy, red the least and

coldest

White is all the colors combined, black the

absence

Infrared Light

Lower energy radiation

Night vision

Microwaves

Low energy radiation

Used to heat food

Also used in

telecommunication

Radio Waves

Have the Lowest

frequencies and highest

wavelengths

Includes FM, AM, and TVs

Radar

REALITY TV

The Particle Description

The Photoelectric effect:

– Emission of electrons when certain light hits

metal

Einstein theorized that light can be modeled as

a photons (particles of light)

– Each photon carries an unique quantum of energy

Maxwell Planck

Theorized that the energy given off by a photon is directly related to the frequency of the radiation emitted.

His equation Energy (E) = hv– h = planck’s constant = 6.626 x 10-24 J•s

Ch. 4.11920s Niels Bohr

Suggested that electrons move

around nuclei in set paths around

the nucleus. (solar system)

7. 1920s Niels Bohr

He said each path is a calculated

energy level

Atom’s electrons can jump to

different energy levels when

absorbing photons

Niels Bohr

States of Atoms

Ground State- An electron’s

lowest energy state or level

Excited State- An electron that is

energized will jump to a higher

energy state or level. Will last for a

short period.

– (resulting light production)

States of Atoms?

Ground State Excited state

Ground State Excited state back to ground

This is how light is produced

And Light production

Niels Bohr

Bohr noticed that different

atoms emitted different

radiation when excited.

Planck’s Flame Test

Spectroscopy

Emission (Line) Spectrum

Atoms and molecules are identified

by these spectrums.

What the Heck is Light?

Ch4.2 Quantum Theory

French Physicist De Broglie

Wave-Particle Duality of electrons

– He proposed that electrons can only exist at certain frequencies, hence the energy they release when excited

Today’s Theory

Werner Heisenberg Uncertainty Principle – It is impossible to determine an

electron’s exact position and speed at the same time.

Along with De Brogile, they disproved Bohr’s “definite orbit” assumption

Schrödinger’s Work

Electrons found in orbitals within different energy levels.

–a calculated region in an atom where there is a high probability of finding electrons.

This led to the Electron Cloud model and Quantum Numbers

Modern Atomic Cloud Model

History Review

Quantum #’s

Principal (n) = number of energy level (1-7)

Angular momentum (l) = sublevels: 0,1,2,or 3

Magnetic (m) = the orbital # (-,0,+)

Spin (s) = electron spin (1/2, -1/2)

Pauli Exclusion Principle

• No two atoms will have the same

configuration or set of quantum #’s

Energy levels

1st level holds 2 e- (s)

2nd level holds 8 e- (s,p)

3rd level holds 8 or 18e- (s,p,d)

4th level holds 18 or 32e- (s, p d, f)

Outer (valence) level holds up to 8 e-

(s, p)

Making Bohr Models

Bohr’s Model

Blue dots represent electrons

Rings represent energy Level,NOT orbit (path)

ELECTRON

CONFIGURATIONS

Ch4.2

Electron Energy Levels, SUBLEVELS, and Orbital's

Electron Configuration

The rows (periods) of the periodic table tell you the energy level

There are 4 sublevels – s,p,d,f

Electron Configuration

Each sublevel contains a different number of orbitals.

S =1, p = 3, d= 5, and f = 7

Orbital is represented by:

Each orbital holds only two electrons with opposite spin.

Writing Atomic Electron Configurations

Three ways of writing configurations.

1. orbital box notation.

2. Electron Configuration (spdf) Notation

3. Noble Gas Configuration

Orbital Notation Rules:

1. Aufbau Principle- electrons are added one

at a time starting at lowest energy level

2. Hund’s Rule - When filling orbitals with in

the same sublevel, fill one electron into each

box before pairing electrons.

Orbital Notation

Arrowsdepictelectronspin

ORBITAL BOX NOTATIONfor He, atomic number = 2

1s

21 s

3. Hund’s Rule1.

2P

S,p,d,f Notation

11 s

value of nsublevel

no. of

electrons

for H, atomic number = 1

Phosphorus

Group 5A

Atomic number = 15

1s2 2s2 2p6 3s2 3p3

[Ne] 3s2 3p3

1s

2s

3s

3p

2p

Aluminum

Group 3A

Atomic # = 13

1s2 2s2 2p6 3s2 3p1

[Ne] 3s2 3p1

All Group 3A elements have [core] ns2 np1

configurations where n is the period number.

1s

2s

3s

3p

2p