CP Chemistry – Chapter 3

Mrs. Albertson

Spring 2001

Lavoisier

Father of Modern Chemistry The first to use truly quantitative research Law of Conservation of Mass Identified components of water as hydrogen

and oxygen http://encarta.msn.com/encnet/refpages/RefAr

ticle.aspx?refid=761571807 http://encarta.msn.com/encnet/refpages/RefM

edia.aspx?refid=461541281

The Structure of the Atom http://www.watertown.k12.wi.us/hs/teachers/

buescher/atomtime.asp http://www.sci.tamucc.edu/pals/morvant/genchem/

atomic/index.htm Atom: Smallest particle of an element that still retains

properties of that element 4th Century B.C. – Democritus first suggested the

idea of atoms Nucleus – contains protons and neutrons Electrons are found outside the nucleus Smaller subatomic particles are not discussed here

Discovery of Electrons

John Dalton Noticed that % of each element in a

compound is always the same: Law of Definite Proportions

Carbon Dioxide Always 27.3% carbon and 72.9 % Oxygen

Dalton’s Atomic Theory

All elements are composed of indivisible particles called atoms

Atoms of the same element are identical; atoms of different elements are different

Atoms of different elements combine in small whole number ratios to form compounds

Chemical reactions occur when atoms are separated, joined, or rearranged

Atoms of one element are not changed into atoms of another element, subdivided, or destroyed

Crooke’s Experiment – 1870’s

Gas Tubes w/2 electrodes (conductors) Anode – positive Cathode – negative

Cathode ray tube – applied voltage & beam of light composed of particles was deflected by a magnet – able to determine they were charged particles

The Discovery of Electrons

J.J. Thomson – Investigating the relationship between matter & electricity

CRT w/fluorescent screen allowed him to measure deflection when a magnet was used

Measure ratio of charge to mass and determined particles were identical regardless of the gas used/subatomic particles

Millikan - 1909

Approximated the mass of an electron to be 1/2000 the mass of an H atom

Current 1/1837 9.109 x 10-31 kg

Protons

Atoms are neutral so a positive charge must exist

Thomson – designed exp to test/H+ moved toward negative end of CRT

Protons identified by 1920 Deflection of + particles varied w/different gases.

Hydrogen had the greatest deflection and smallest mass

Mass of Proton 1.673 x 10-27 kg

Thomson’s “Plum Pudding” Model

Nobel Prize 1907 Pudding was + charge and most of the mass of the

atom Plums: - charged electrons spread throughout to

make the atom neutral Ions: + / - charged atoms: result from the loss or

gain of electrons Cations – positive charge / lost electrons Anions – negative charge / gain electrons http://www.sci.tamucc.edu/pals/morvant/genchem/atomic/

page6.htm

Radioactivity Discovered

1896 – Radioactivity discovered in Uranium by Becquerel Radiation: energy that is emitted from a

source and travels through space Radioactivity: spontaneous radiation from the

nucleus of an atom Marie/Pierre Curie – radium & polonium

Radioactivity

By 1900 3 types of radiation identified Alpha – He ions w no elctrons; 1/10th the

speed of light; stopped by paper or clothing Beta – electrons at high speeds / stopped by

a few mm of Al Gamma – form of electromagnetic radiation;

more energetic than x-rays; stopped by several cm of Pb or more concrete/ no mass or charge

Rutherford’s Gold Foil

Resulted in a new model of the atom Atoms contain a small dense nucleus Electrons move around like bees in a hive Diameter of nucleus 1/100,000 the size of the atom 1920 Rutherford proposed neutral particles with the

same mass as protons http://micro.magnet.fsu.edu/electromag/java/

rutherford/ http://www.brainpop.com/science/matter/

atomicmodel/index.weml?&tried_cookie=true

Chadwick

Credited with the discovery of neutrons Nobel Prize – 1935 Neutron Mass – 1.675 x 10-27 kg

Forces in the Nucleus

Like charges normally repel Protons are strongly attracted to one

another in the nucleus Also neutron/neutron and neutron/proton

attractions These are the result of

NUCLEAR FORCES

Atomic Number & Mass Number

Atomic number – the number of protons in the nucleus; defines what element an atom is

Mass number Protons + Neutrons = Mass Number

Amu – atomic mass units – 1/12 the mass of a carbon-12 atom

1 proton = 1.007276 amu 1 neutron = 1.008665 amu

http://www.sci.tamucc.edu/pals/morvant/genchem/atomic/page8.htm

Isotopes

Atoms of the same element with different numbers of neutrons http://www.sci.tamucc.edu/pals/morvant/genchem/atomic/

page9.htm Nuclide – general term for any isotope of any element Each isotope has a % abundance in nature Symbols for isotopes:

Lithium – 6 / Lithium – 7 Isotopes differ by

Number of neutrons Mass number Atomic mass

Isotopes Cont.

Of 1500 known isotopes, only 264 are stable; others are radioactive

Radioactive decay – alpha or beta particles are emitted and the nucleus changes to form a new element or isotope – continues until a stable form is reached

Isotopes Cont

Atomic mass – average of the mass of an elements isotopes based on % abundance Carbon – 12.011 amu

Carbon 13 = 1.11 % Carbon 12 = 98.89 %

Example Problem – Find average atomic mass of carbon:

.0111 x 13amu = .1443 amu .9889 x 12 amu = 11.8668 amu

12.011 amu Take relative abundance x mass of isotope and add together

Example Problem

Using the following information, determine the atomic mass of chlorine: Two isotopes are know: chlorine-35 (mass=35.0 amu) and chlorine-37 (mass=37.0 amu). Their relative abundances are 75.4% and 24.6% respectively.

Sample ProblemsElement P N E Mass

#

*Avg.

Atomic mass

Nitrogen-15

6 7

8 15

Mn(+2) 30

Relating Mass to Number of Atoms

The MOLE SI unit for amount of substance Amt. of substance that contains as many

particles as there are atoms in exactly 12 g of carbon-12

Counting unit – just like a dozen Avogadro’s Number

Experimentally determined: 6.022 x 1023

Particles & the MOLE

4 Types of Particles Atoms Ions Molecules Formula Units

There are 6.022 x 1023 particles in 1 mole of any pure substance

What is a pure substance?

Classification of Matter

C la ss if ica tion o f M atte r

C a rb o n , N itrog en

E lem en ts

N a C l, H (2 )O

C o m p o un ds

P u re S u b sta nces

S a lt W a te rA ny S o lu tion

H o m o ge n eo us

B e ef S te w , S an d,D iffe re n t P ha ses

H e te rog en eo us

M ix tu res

M atte r

Molar Mass

The mass of 1 mole of a pure substance in a unit of g/mol

Equal to atomic mass A molar mass of an element contains 1 mol of atoms Examples:

Iron (Fe) = 55.85 g/mol55.85 g of iron contains 6.022 x 1023 atoms

Water (HOH) = 2 mol H atoms x 1.01 g/mol = 2.02 g + 1 mol O atoms x 16.00 g/mol = 16.00 g

18.02 g

Gram/Mole Conversions

What is the mass of 5.00 mol of Ni (in grams)? Examine Plan Organize Evaluate

How many moles are in 70 g of carbon?

Conversions w/Avogadro’s Number

How many atoms are in 3.0 mol of He? How many atoms are in 52.63 g of Na? How many H atoms are in 3 mol of water?

REMEMBER IF GOING FROM A COUNTING UNIT TO GRAMS OR GRAMS TO COUNTING UNIT – GO TO MOLES FIRST

gmolatoms atomsmolg