Chemistry Unit 3 Atomic Structure (Ch.3). 3-1 Early Models of Atoms zDemocritus (450 BC) zProposed...

transcript

Chemistry Unit 3

Atomic Structure (Ch.3)

3-1 Early Models of AtomsDemocritus (450

BC)Proposed that all

matter was made of tiny indivisible particles.

He called these particles atomos (meaning indivisible).

We call them atoms.Good looking guy!

An atom is the smallest particle of an element that retains the identity of that element.

If we repeatedly cut a piece of Al, the smallest possible piece is an atom of Al.

Classic model of an atom

Aristotle

Didn’t agree with Democritus.

Believed matter was continuous and made up of only one substance called “hyle”

It wasn’t until the 1700’s when his ideas were reexamined.

Newton and Boyle (1600s)

Published articles stating their belief in the atomic nature of elements

They had no proof

Antoine Lavoisier (1770’s)

French, The “Father of Modern Chemistry”

Discovered the law of conservation of matter.

Matter is neither created nor destroyed.

Joseph Proust (1799)

French ChemistDeveloped The

Law of Definite Proportions

Compounds always contain elements in the same proportion by mass.

Law of Definite Proportions

H20 (by mass is always)

88.9% Oxygen, 11.1% Hydrogen

If we had an 80g sample of H20 how much is O?

.889 x 80 = 71gHow much is H?.111 x 80 = 9g

John Dalton (1803)

Proposed the atomic theory of matter, which is the basis for present atomic theory

John Dalton,

English schoolteacher

Atomic Theory of Matter

Each element is composed of extremely small particles called atoms.

All atoms of a given element are identical, but differ from those of any other element.

Which element is this?

Atomic theory of matter

When elements unite to form compounds, they do so in a ratio of small whole numbers. This is called the Law of Multiple Proportions.

Ex: C and O can combine to form CO or CO2, but not CO1.8.

Dalton’s Model of an Atom

All matter is composed of tiny particles

J.L. Gay-Lussac (early 1800s)

Observed that working with gas reactions at constant volume, temperature and pressure are directly related.

He named the discovery of this relationship Charles Law, which is represented by P1/T1=P2/T2.

Amadeo Avogadro (early 1800s) – Italian Physicist

Explained Gay-Lussac’s work using Dalton’s theory: Equal volumes of gases at the same temp/pres have the same number of gas molecules.

Michael Faraday (1839)Suggested that

atomic structure was related to electricity.

Atoms contain particles that have electrical charges.

Positive (+)Negative (-)Opposite charges

attractLike charges repel

William Crookes (1870’s)

English Physicist

Developed the cathode ray tube to find evidence for the existence of particles within the atom.

J.J. Thomson (1896)Used a cathode

ray tube (CRT) to identify negatively charged particles, called electrons.

Determined the ratio of an electron’s charge to its mass.

Developed the “plum pudding” model of an atom.

Cathode ray bending toward a positive charge

Plum Pudding Model

Atoms are composed of randomly arranged charged particles

Robert Millikan (early 1900s)

US PhysicistUsed the oil drop

experiment to prove the electron has a negative charge

Was able to determine the charge of the electron

Millikan’s Oil Drop Experiment

Bothe/Chadwick (early 1930s)

English

Found high energy particles with no charge with the same mass as the proton called neutrons.

Ernest Rutherford (1909)Used the gold foil

experiment to prove the atom is mostly empty space.

Rutherford concluded that all of an atom’s positive charge, and most of its mass is located in the center, called the nucleus. Analogy: thumb nail and the 50 yard line.

98% of the particles passes straight through 2% of the particles deflected off at varying angles 0.01% of the particles bounced straight back

Rutherford’s PlanetaryModel of an atom

Positive charge and majority of mass located in the nucleus.

Negatively charged electrons orbit the nucleus like planets. Most of an atom is empty space!

Problem

He thought a moving electrical charge (-) in a curved path should lose energy (give off light).

If it did, it would fall into the (+) nucleus.

Why don’t the (-) electrons fall into the (+) nucleus?

Atom:The smallest particle of an element that has the properties of that element. Make up of nucleus

consists of protons and neutrons

Surrounded by an electron cloud

Electron cloud

Sub-Atomic Particles

Protons Positively (+) charged The number of protons

in an atom refers to the atomic number (Z)

Composed of 3 quarks (2 up, 1 down)

Mass= 1.6726 x 10-

27kg Atomic mass 1

amu (µ)

Neutrons Found in nucleus Neutral (no) charge composed of 3

quarks (1 up, 2 down)

Atomic mass 1 amu (µ)

Isotopes- atoms of the same element that have a different number of neutrons.

Electrons Found in electron

clouds surrounding the nucleus.

Negative (-) charge Mass = .00091

x 10-27 kg 1800 times smaller

than protons & neutrons

Mass 0 amu (µ)

Sub-atomic particles

Electrons Orbit the nucleus at

very high speed in energy levels (electron clouds).

Negatively (-) charged

Have no mass (when compared to protons and neutrons)

Atomic Number = Protons

The atomic number of an element is the number of protons an element has.

Located above the symbol of the element

The number of protons determines the identity of the element.

Each element has a different atomic number

Neils Bohr (1913)

Improved Rutherford’s work by saying electrons do not lose energy in the atoms so they will stay in orbit

Stated there are definite levels in which the electrons follow set paths without gaining or losing energy (Planetary Model)

Each level has a certain amount of energy associated with it and the electrons can only jump levels if they gain or lose energy

Could not explain why (-) electrons don’t fall into the (+) nucleus.

Energy Levels In the ground state

for an atom, electrons are at their lowest, most stable energy levels.

In the excited state, atoms require energy and electrons move to a higher energy level.

How many electrons are in an atom?

For an atom to have an overall neutral charge the number of electrons must equal the number of protons.

#Protons=#electrons

What element is this?

Mass number

The Mass number of an atom is the sum of the mass of protons and neutrons

Located below the symbol of the element

Atomic mass is measured in amu’s, (atomic mass units)

Based on Carbon having a mass of 12

Mass = Protons + Neutrons

How many neutrons are in an atom?

Mass=Protons+Neutrons 195= 78 + Neutrons 195-78= Neutrons Platinum has 117

Neutrons Find the number of

neutrons in: Hydrogen Carbon Helium

Potassium Boron Gold

Mass =Protons + Neutrons Hydrogen (H) 1 =1 + Neutrons

Hydrogen has 0 neutrons Helium (He) 4 = 2 + Neutrons

Helium has 2 neutrons Boron (B) 11 = 5 + Neutrons

Boron has 6 neutrons Carbon (C) 12 = 6 + Neutrons

Carbon has 6 neutrons Potassium (K) 39 = 19 + N

Potassium has 20 neutrons Gold (Au) 197 = 79 + N

Gold has 118 neutrons

Atomic MassThe average

mass of all of the isotopes of an element.

Aka: average atomic mass number, or atomic weight.

Isotopes:Atoms of the same element with different masses.

Average Atomic Mass

Ne-20 has a mass of 19.992 amu (u), and Ne-22 has a mass of 21.991 amu (u). In any sample of 100 Ne atoms, 90 will be Ne-20. Calculate the average atomic mass of Ne.

.90 x 19.992 = 17.9928.10 x 21.991 = 2.1991avg mass = 20.1919 amu

IonsAn atom that has

gained or lost an electron.

It acquires a net electrical charge.

If an atom loses an electron (oxidation) it has more protons than electrons and has a net positive charge. (cation)

10 e-Na+

If an atom gains an electron (reduction) it has more electrons than protons and has a net negative charge.(anion)

Full octet7 valence e-

Ionic ChargesCharge of ion = # protons - #

electronsWhat is the charge of a magnesium

atom that loses 2 electrons?Number of protons 12-Number of electrons 10charge of ion +2Mg2+ or Mg+2

Charge is written to the upper right of the symbol.

Representations of atoms

General form: (Elemental Notation)

X = Element SymbolA = Atomic Mass

(P + N)Z = Atomic Number

(P)Ionic Charge

Charge

What is the atomic structure?

Determine the number of:

P =N =e =

P = 11N = 12e-= 10

P =N = e- =

P = 53N = 74e- = 54

Put into elemental notation

Atomic # = 29Atomic Mass = 64Ionic charge = +2 ?

How many electrons?

Atomic # = 29Atomic Mass = 64Ionic charge = +2# of electrons =

37 Protons48 Neutrons36 Electrons

37 Protons48 Neutrons36 Electrons Rb

Max Planck (early 1900s)

Proposed Planck’s Theory which says that energy is given off in little packets or particles called quanta which is based on the particle nature of light

Each quantum of energy corresponds to the different energy levels for electrons.

Proposed the equation: E=hf, where E is energy, f is frequency, and h is Planck’s constant (6.63 x 10^-34 J/Hz)

De Broglie (1923)

Suggested that if waves can have a particle nature, particles can have a wave nature, known as the “wave-particle duality” principle

Wondered why the positive nucleus and negative electrons do not attract. Proposed that electrons moved so fast (speed of light) that they had properties of waves instead of particles.

The Study of WavesWave: a progressive disturbance

propagated from point to point in a medium or space without progress or advance by the points themselves

Types of Waves

Mechanical: a wave that requires an energy source and an elastic material medium to travel.

Electromagnetic: a wave that does not require a material medium to travel; it propagates by electric and magnetic fields

Wave Travel

Transverse: displacement of the medium is perpendicular to the direction of propagation of the wave.

Longitudinal: displacement of the medium is parallel to the direction of propagation of the wave

Properties of Waves

Wavelength (ג): The distance between any part of the wave (peak) and the nearest part that is in phase with it (another peak). Standard unit is meters (m).

Frequency (f ): The number of peaks which pass a given point each second. Standard unit is cycles per second which is a hertz (Hz).

Amplitude (A): The maximum displacement of a vibrating particle from its equilibrium position. Standard unit is meters (m).

Velocity (v): the distance a wave (peak) travels in a given time. Standard unit is meters per second (m/s).

Energy (E): The energy of a single photon of radiation of a given frequency. Standard unit is the joule (J).

Some relationships between the properties of waves are represented by the equations:

V=f*ג and E=h*f , where h=6.63x10^-34 J/Hz

Werner Heisenberg (1927)

Proposed his “Heisenberg uncertainty principle”, which says that the position and momentum of an electron cannot simultaneously be measured and known exactly.

The arrangement of electrons is discussed in terms of the probability of finding an electron in a certain location.

Erwin Schrodinger (1926)

Studied the wave nature of the electron and developed mathematical equations to describe their wave-like behavior.

The most probable location of the electrons can be found and the plot of this probability is called the charge cloud model.

The four quantum numbers

Principal Quantum Number (n) Refers to the energy

levels in the atom which is the distance from the nucleus and designated with a positive whole number (n=1,2,3,etc)

Wavelength of emitted photon is determined by the “energy jump” between energy levels

Energy levels (or shells) means electrons are contained in an area where the probability of finding the electron is 90%

Angular Momentum Quantum Number (l ) Refers to the sublevel

(within an energy level) which is one or more “partitions” each with a slightly different energy.

The number of sublevels in a particular energy level is equal to the principal quantum number (n).

The types of sublevels include: s, p, d, f, etc.

The four quantum numbers (continued)

Magnetic Quantum Number (m)

Refers to the orientation in space of the electrons in a sublevel

Can have any whole number value from -1 to +1 which will tell how many orbitals are in a sublevel.

A maximum of 2 electrons per orbital.

Sublevel # of Orbitals Total # of electronsspdf

Four Quantum Numbers (continued)

Spin Quantum Number (s) Refers to the spin of the electron. Pauli Exclusion Principle : if two

electrons occupy the same orbital, they must have opposite spin. Half-filled orbital: _____Filled orbital: _____

Permissible Values of Quantum Numbers for Atomic Orbitals

n l m Orbital # of Subshells #of Orbitals Max # of Electrons1 0 0 1s 1 1 2

2 0 0 2s 2 1 2 1 -1,0,1 2p 3 6

3 0 0 3s 3 1 2 1 -1,0,1 3p 3 6 2 -2,-1,0,1,2 3d 5 10

4 0 0 4s 4 1 2 1 -1,0,1 4p 3 6 2 -2,-1,0,1,2 4d 5 10

3 -3,-2,-1,0,1,2,3 4f 7 14

Distribution of Electrons for Different Elements (Electron Configuration)

Electrons will occupy the lowest energy levels and sublevels first.

Notation:

Type of Orbital (sublevel)

Principal Quantum Number, n (energy level)

Number of electrons

Orientation of Orbital

Long Notation: Pyramid Filling

“Rule of thumb” for filling electrons at the lowest energy level possible.

Give the long notation electron configuration for:

Give the short notation

Orbital Diagrams

Usually only done for the outer shell electrons, which always includes the s and p orbitals.

Electron Dot DiagramsShows the outer shell electrons

for elements.

Chemistry Unit 3 Atomic Structure (Ch.3). 3-1 Early Models of Atoms zDemocritus (450 BC) zProposed...

Documents