Atomic theory
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Atomic Theory. Defining the Atom. The Greek philosopher Democritus (460 B.C. – 370 B.C .) was among the first to suggest the existence of atoms (from the Greek word “atomos”) He believed that atoms were indivisible and indestructible

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Atomic theory

Atomic Theory


Defining the atom

Defining the Atom

  • The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”)

    • He believed that atoms were indivisibleand indestructible

    • His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – but just philosophy


Dalton s atomic theory experiment based

Dalton’s Atomic Theory (experiment based!)

All elements are composed of tiny indivisible particles called atoms

Atoms of the same element are identical. Atoms of any one element are different from those of any other element.

John Dalton

(1766 – 1844)

Atoms of different elements combine in simple whole-number ratios to form chemical compounds

In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.


Dalton s atomic theory 1808

Dalton’s Atomic Theory (1808)

  • Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements.

  • Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same.

  • Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions.

2.1


Atomic theory

2

Law of Multiple Proportions

2.1


Atomic theory

16 X

+

8 Y

8 X2Y

Law of Conservation of Mass

2.1


Structure of the nuclear atom

Structure of the Nuclear Atom

Dalton Wasn’t Exactly Correct…..

  • One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles:

    • Electrons, protons, and neutrons are examples of these fundamental particles

    • There are many other types of particles, but we will study these three


Discovery of the electron

Discovery of the Electron

In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron


Atomic theory

J.J. Thomson, measured mass/charge of e-

(1906 Nobel Prize in Physics)

2.2


Atomic theory

Cathode Ray Tube

2.2


Thomson s atomic model

Thomson’s Atomic Model

J. J. Thomson

Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.


Modern c athode r ay t ubes

Modern Cathode Ray Tubes

Television

Computer Monitor

  • Cathode ray tubes pass electricity (electrons) through a gas that is contained at a very low pressure.


Mass of the electron

Mass of the Electron

Mass of the electron is

9.11 x 10-28 g

The oil drop apparatus

1916 – Robert Millikan determines the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge


Atomic theory

Measured mass of e-

(1923 Nobel Prize in Physics)

e-charge = -1.60 x 10-19 C

Thomson’s charge/mass of e- = -1.76 x 108 C/g

e- mass = 9.10 x 10-28 g

2.2


Atomic theory

(Uranium compound)

2.2


Atomic theory

2.2


Atomic theory

(1908 Nobel Prize in Chemistry)

  • particle velocity ~ 1.4 x 107 m/s

    (~5% speed of light)

  • atoms positive charge is concentrated in the nucleus

  • proton (p) has opposite (+) charge of electron (-)

  • mass of p is 1840 x mass of e- (1.67 x 10-24 g)

2.2


Rutherford s findings

Rutherford’s Findings

  • Most of the particles passed right through

  • A few particles were deflected

  • VERY FEW were greatly deflected

“Like howitzer shells bouncing off of tissue paper!”

Conclusions:

The nucleus is small

The nucleus is dense

The nucleus is positively charged


Atomic theory

Rutherford’s Model of

the Atom

atomic radius ~ 100 pm = 1 x 10-10 m

nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m

“If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”

2.2


The rutherford atomic model

The Rutherford Atomic Model

  • Based on his experimental evidence:

    • The atom is mostly empty space

    • All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “nucleus”

    • The nucleus is composed of protons and neutrons (they make the nucleus!)

    • The electrons distributed around the nucleus, and occupy most of the volume

    • His model was called a “nuclear model”


Chadwick s experiment 1932

Chadwick’s Experiment (1932)

H atoms - 1 p; He atoms - 2 p

mass He/mass H should = 2

measured mass He/mass H = 4

· In 1932 James proved the existence of neutral particles in an

atom · James said that the neutrons were just about the same weight as

protons · He discovered this by using alpha rays, which are charged, and

therefore repelled by considerable electrical forces present in the

nuclei of heavier atoms · Chadwick led the way to the starting of penetrating and splitting

the nuclei of atoms. · Also led the way to the fission of uranium 235, which eventually

created the atomic bomb


Homework

HOMEWORK

1. Describe JJ Thompson’s CRT (Cathode Ray Tube) experiment & how it showed that atoms contain particles he called “electrons.”

2. Describe JJ’s model of the atom.

3. Explain Rutherford’s scattering experiment and what it helped to prove. Also, how did it disprove Thompson’s model?

  • Describe Rutherford’s atomic model.

  • What led to Chadwick’s discovery.


Atomic theory

mass p+ = mass no = 1840 x mass e-

Mass (amu)

0

1

1

2.2


Atomic theory

A

X

Mass Number

Element Symbol

Z

Atomic Number

2

3

1

H (D)

H (T)

H

1

1

1

235

238

U

U

92

92

Atomic number (Z) = number of protons in nucleus

Mass number (A) = number of protons + number of neutrons

= atomic number (Z) + number of neutrons

Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei

2.3


Atomic theory

2.3


Atomic theory

14

11

C

C

6

6

How many protons, neutrons, and electrons are in

How many protons, neutrons, and electrons are in

?

?

Do You Understand Isotopes?

6 protons, 8 (14 - 6) neutrons, 6 electrons

6 protons, 5 (11 - 6) neutrons, 6 electrons

2.3


Atomic theory

Ions: Atoms with different number of protons (p+) and electrons (e-)

11 protons

11 electrons

11 protons

10 electrons

Na+

Na

17 protons

18 electrons

17 protons

17 electrons

Cl-

Cl

cation – ion with a positive charge

If a neutral atom loses one or more electrons

it becomes a cation.

anion – ion with a negative charge

If a neutral atom gains one or more electrons

it becomes an anion.


Atomic theory

How many protons and electrons are in ?

27

3+

Al

13

Do You Understand Ions?

13 protons, 10 (13 – 3) electrons

78

2-

How many protons and electrons are in ?

Se

34

34 protons, 36 (34 + 2) electrons

2.5


Atomic theory

Practice:

  • Write the elemental symbol if 10 p+, 11no, 10e-

21Ne

2) Write the elemental symbol if 20p+, 20no, 18e-

40Ca2+

2.5


Atomic theory

Noble Gas

Halogen

Alkaline Earth Metal

Period

Alkali Metal

Group

2.4


Atomic theory

Chemistry In Action

Natural abundance of elements in Earth’s crust

Natural abundance of elements in human body

2.4


Atomic theory

H2

H2O

NH3

CH4

A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds

A diatomic molecule contains only two atoms

H2, N2, O2, Br2, HCl, CO

A polyatomic molecule contains more than two atoms

O3, H2O, NH3, CH4

2.5


Atomic theory

11 protons

11 electrons

11 protons

10 electrons

Na+

Na

17 protons

18 electrons

17 protons

17 electrons

Cl-

Cl

An ion is an atom, or group of atoms, that has a net positive or negative charge.

cation – ion with a positive charge

If a neutral atom loses one or more electrons

it becomes a cation.

anion – ion with a negative charge

If a neutral atom gains one or more electrons

it becomes an anion.

2.5


Atomic theory

A monatomic ion contains only one atom

Na+, Cl-, Ca2+, O2-, Al3+, N3-

A polyatomic ion contains more than one atom

OH-, CN-, NH4+, NO3-

2.5


Atomic theory

2.5


Atomic theory

2.6


Atomic theory

molecular

empirical

H2O

A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance

An empirical formula shows the simplest

whole-number ratio of the atoms in a substance

H2O

CH2O

C6H12O6

O3

O

N2H4

NH2

2.6


Atomic theory

  • ionic compounds consist of a combination of cations and an anions

  • the formula is always the same as the empirical formula

  • the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero

The ionic compound NaCl

2.6


Atomic theory

1 x +2 = +2

1 x +2 = +2

2 x +3 = +6

1 x -2 = -2

3 x -2 = -6

2 x -1 = -2

Formula of Ionic Compounds

Al2O3

Al3+

O2-

CaBr2

Ca2+

Br-

Na2CO3

Na+

CO32-

2.6


Atomic theory

2.6


Atomic theory

2.7


Polyatomic ions

Polyatomic Ions


Writing formulae diatomic polyatomic

WRITING FORMULAE (DIATOMIC & POLYATOMIC)

Polyatomic

1) Potassium Nitrate

2) Sodium Sulfate

3) Potassium Dichromate

4) Ammonium Phosphate

5) Copper I Carbonate

6) Iron (III) Cyanide

7) Silver Sulfite

8) Tin (II) Nitrite

9) Calcium Hydroxide

10) Boron Acetate

Mixed

11. Sodium Sulfide

12. Sodium Bicarbonate

13. Potassium Oxide

14. Potassium Permanganate

15. Magnesium Chloride

16. Calcium Chlorate

17. Iron (II) Phosphate

18. Copper Nitride

19. Ammonium Phosphide

20. Aluminum Acetate


Chemical nomenclature

Chemical Nomenclature

  • Ionic Compounds

    • often a metal + nonmetal

    • anion (nonmetal), add “ide” to element name

barium chloride

BaCl2

potassium oxide

K2O

magnesium hydroxide

Mg(OH)2

potassium nitrate

KNO3

2.7


Atomic theory

  • Transition metal ionic compounds

    • indicate charge on metal with Roman numerals

iron(II) chloride

FeCl2

2 Cl- -2 so Fe is +2

FeCl3

3 Cl- -3 so Fe is +3

iron(III) chloride

Cr2S3

3 S-2 -6 so Cr is +3 (6/2)

chromium(III) sulfide

2.7


Atomic theory

Hydrates:

  • ionic compounds with Water Molecules contained in their crystalline structure.

CuCl2 · 2 H2O

Copper Chloride Dihydrate

Fe(SO3) · 5 H2O

Iron (II) Sulfate Pentahydrate

Cu(NO3)2 · 4 H2O

Copper Nitrate Tetrahydrate

2.7


Atomic theory

Hydrates:

  • USE PREFIXES TO DENOTE THE # OF WATER MOLECULES

1 =

mono

6 =

hexa

2 =

di

7 =

hepta

3 =

tri

8 =

octa

4 =

tetra

9 =

nona

5 =

penta

10 =

deca


Atomic theory

  • Molecular compounds

    • nonmetals or nonmetals + metalloids

    • common names

      • H2O, NH3, CH4, C60

    • element further left in periodic table is 1st

    • element closest to bottom of group is 1st

    • if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom

    • last element ends in ide

2.7


Atomic theory

TOXIC!

Laughing Gas

Molecular Compounds

HI

hydrogen iodide

NF3

nitrogen trifluoride

SO2

sulfur dioxide

N2Cl4

dinitrogen tetrachloride

NO2

nitrogen dioxide

N2O

dinitrogen monoxide

2.7


Atomic theory

2.7


Atomic theory

nitric acid

HNO3

carbonic acid

H2CO3

H2SO4

sulfuric acid

An acid can be defined as a substance that yields

hydrogen ions (H+) when dissolved in water.

  • HCl

    • Pure substance, hydrogen chloride

    • Dissolved in water (H+ Cl-), hydrochloric acid

An oxoacid is an acid that contains hydrogen, oxygen, and another element.

HNO3

2.7


Atomic theory

2.7


Atomic theory

2.7


Atomic theory

2.7


Atomic theory

sodium hydroxide

NaOH

potassium hydroxide

KOH

Ba(OH)2

barium hydroxide

A base can be defined as a substance that yields

hydroxide ions (OH-) when dissolved in water.

2.7


Atomic theory

2.7


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