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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