Chapter 12: DNA and RNA Ferguson 2014 Honors Biology/Chemistry.
Honors Chemistry chapter 3
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Transcript of Honors Chemistry chapter 3
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HONORS CHEMISTRY CHAPTER 3
Atoms: The Building Blocks of Matter
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ICONS OFEARLY ATOMIC
THEORY
PART 1
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ICONS IN EARLY ATOMIC THEORY Democritus [400 B.C]
Greek philosopherHypothesized: Nature
has a basic indivisible particle of which everything is made of Called this particle an
atomGreek “atomos” =
indivisible
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1790S – DISCOVERY OF BASIC LAWS Law of Conservation of Mass
Mass is neither created nor destroyed during ordinary chemical reactions or physical changes
Law of Definite Proportions A chemical compound contains the same
elements in exactly the same proportions by mass regardless of size of sample or source of compound i.e. Every sample of table salt is made of 39.34%
Na and 60.66% Cl i.e. H2O always has 2 atoms of H and 1 atom of
O
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BASIC LAWS CONTINUED Law of Multiple Proportions
If two or more different compounds are composed of the same two elements then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers i.e. CO and CO2
CO = 1.00g of C and 1.33 g of OCO2 = 1.00 g of C and 2.66 g of OThe ratio of the second element is 2.66 to 1.33 or 2
to 1
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ICONS OFEARLY ATOMIC THEORY CONTINUED
John Dalton [1808]English schoolteacher – liked nature and
weatherDeveloped: Dalton’s Atomic Theory
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DALTON’S ATOMIC THEORY1. All matter is composed of extremely
small particles called atoms2. Atoms of a given element are identical
in size, mass and other properties and are different from atoms of other elements
3. Atoms cannot be subdivided, created, or destroyed
4. Atoms of different elements combine in simple whole number ratios to form chemical compounds
5. In chemical reactions, atoms are combined, separated or rearranged
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ISSUES WITHDALTON’S ATOMIC THEORY
Atoms can be split into even smaller particles (nuclear chemistry) and aren’t indivisible i.e. nucleus, protons, electrons
A given element can have different masses i.e. isotopes
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STRUCTURE OF THE ATOM Today’s definition of the atom
Atom = Smallest particle of an element that retains the chemical properties of that element Two regions
Nucleus Very dense, small center of the
atoms Protons and neutrons
Electron Cloud Region occupied by electrons
Subatomic particlesProtons, neutrons, electrons
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ICONS OF EARLY ATOMIC THEORY CONTINUED
J.J. Thomson [1897]Discovered: The 1st
subatomic particle: the negatively charged electron
Used a Cathode Ray Experiment Cathode Ray Tube –
Electric current passed through a metal disk to another metal disk in a gas at low pressure (vacuum sealed tube)
i. e. neon signs and ‘old-fashioned’ television sets
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CATHODE RAY EXPERIMENT When a current passed
through the cathode ray tube, the surface of the tube opposite the cathode glowed Glow was hypothesized to be
stream of particles called a cathode ray
Ray affected by magnetic fields Attracted to positive charge Deflected from negative
charge http://www.youtube.com/watch?v
=7YHwMWcxeX8&NR=1
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DISCOVERY OF THE1ST SUBATOMIC PARTICLE Thomson measured the ratio of the
charge of the particles to their massSame ratio no matter what metal or gas
was usedNamed this particle an electron
http://www.youtube.com/watch?v=IdTxGJjA4Jw&feature=related
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THOMSON’SPLUM PUDDING MODEL Atoms are electrically neutral
Must have positive charges to balance the negatively charged electrons
Electrons have a lot less mass than atomsOther particles must account for their mass
Plum Pudding Modelpositively charged sphere with electrons
dispersed through it
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ICONS OFEARLY ATOMIC THEORY CONTINUED
Robert Millikan [1909]Discovered: The
measurement of an electron charge
Oil Drop Experiment Measured the
difference in velocity of oil dropletsCharged droplets
(ionizing radiation) vs. uncharged
http://www.youtube.com/watch?v=XMfYHag7Liw&feature=related
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ICONS OFEARLY ATOMIC THEORY CONTINUED Ernest Rutherford (with Hans Geiger and
Ernest Marsden) [1911] Discovered: A new atomic model Gold Foil Experiment
Bombarded thin piece of gold foil with alpha particles Expected alpha particles to pass through with minimal
deflection Surprised when 1 in 8000 deflected back to source
It was “as if you had fired a 15 inch artillery shell at a piece of tissue paper and it came back and hit you”
http://www.youtube.com/watch?v=wzALbzTdnc8&feature=related
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RUTHERFORD’SNEW MODEL OF THE ATOM Discovered the nucleus is a small
densely packed volume of positive chargeSize comparison
Nucleus = marble Whole Atom = football field
At this point in history, we were not sure where the electrons were – stay tuned for more in Chapter 4
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INSIDE THE ATOMPART 2
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INSIDE THE NUCLEUS 2 types of particles
Protons positively charged = +1 made up of quarks
Neutrons neutral = 0 charge Made up of quarks
Mass in the nucleus Protons = 1.673 x 10-27
Neutrons = 1.675 x 10-27
To simplify, both have mass of 1 amu (atomic mass unit)
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HOW DOES THE NUCLEUS STAY TOGETHER?
Strong Nuclear Forces Two protons extremely
close = strong attraction Two neutrons extremely
close = strong attraction Neutrons and Protons extremely close = strong attraction
Strong nuclear forces overcome the repulsion of like positive charges to keep the nucleus together!!!
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WHERE ARE THE ELECTRONS? In the Electron Cloud
A cloud of negative charge outside of the nucleus
More on this later........ Electrons = Negatively charged particles
with almost no mass (9.109 x 10-31)
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CHARACTERISTICS OF ATOMS Atomic Number
Equal to the number of protons and specific to each type of element
Identifies the element # of protons is what give that element its
characteristic properties Elements with different protons are NOT THE SAME
ELEMENT!!!
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NEUTRAL ATOMS Neutral atoms
total positive charge equals the total negative charge # protons (+1 each) = # electrons (-1
each)
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ISOTOPES Atoms of the same element (i.e.
same # of protons) that have differing number of neutrons
Isotopes of the same elementhave different massesdo not differ significantly in chemical
behavior
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MASS NUMBER Mass number = #protons + #
neutrons
Average Atomic Mass Every element has isotopes The periodic table takes into account
all naturally occurring isotopes of an element and averages them
Element
Atomic Number
# of Protons
# of Neutrons
Mass Number
Carbon 6 68 16
Nitrogen
15
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PROPERTIES OF SUBATOMIC PARTICLESParticle Symbol Charge Mass
NumberElectron e-, 0e -1 0
Proton p+, 1H +1 1
Neutron n◦, 1n 0 1
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IONS Atoms with a charge
Negative – more electrons than protonsPositive – more protons than electrons
Charge = #protons - # electrons Magnesium atom with 12 protons and
10 electrons has a charge of +2
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AVERAGE ATOMIC MASS Average Atomic Mass listed on the
periodic table UNIT is amu = atomic mass unit
1 amu is a standard Equal to 1/12 the mass of a C-12 atom
Takes into account all an elements isotopes and the frequency of each isotopes occurrence in natureHow to Calculate Average Atomic Mass
Mass of isotope #1 x abundance
in nature (decimal)
+ Mass of isotope #2 x abundance
in nature (decimal)
+ … =
Average Atomic Mass
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EXAMPLE OF CALCULATING THE AVERAGE ATOMIC MASS – HYDROGEN
There are two naturally occurring isotopes of hydrogenHydrogen with 1 proton and zero neutronsHydrogen with 1 proton and one neutron
Differentiating between the two isotopes (symbol – mass number)
Calculation:
Hydrogen Isotopes
Element – mass #
Atomic Mass
Naturally occurring
abundance %1 proton
+ 0 neutronsH-1 1.007825
amu99.9885
1 proton+ 1 neutron
H-2 2.014102 amu
0.0115
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PART 3: THE MOLE
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AMADEO AVOGADRO Amadeo Avogadro [1776]
Lawyer turned professor of mathematical physics
Theorized: equal volumes of all gases at the same temperature and pressure contain the same number of particles.
After Avogadro’s death Avogadro’s number was determined
Avogadro’s number is simply a unit of measure 1 mole = 6.023 x 1023 of any substance
Typically used to talk about particles (atoms, compounds, etc.)
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THE MOLE UNITPUT INTO PERSPECTIVE!!! One mole of rice grains is more grains
than the total number of grains grown since the beginning of time.
A mole of rice would occupy a cube about 120 miles on each edge.
A mole of marshmallows would cover the US to a depth of 600 miles
A mole of hockey pucks would be equal in volume to the moon
A mole of basketballs would just about fit perfectly into a ball bag the size of the earth.
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MOLE VIDEO http://www.youtube.com/watch?v=Hj83
oRHdezc
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MOLE CALCULATIONS 1 mole = 6.02 x 1023 of anything (atoms,
molecules, formula units, particles, etc.) Use dimensional analysis when solving:
Conversion factor:1 mole = 6.02 x 1023 atoms, particles, formula units, etc.
Practice:A. If I have 3.5 moles of carbon atoms, how many
molecules do I have?
B. If I have 5.43 x 1031 molecules of carbon dioxide, how many moles do I have?
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MOLAR MASS Molar Mass
The mass of one mole of a substance The molar mass of an element canbe found on the periodic table
Same as the average atomic mass1 amu = 1 gram/mole
E.g. Average atomic mass of C = 12.011 amus Molar mass of C = 12.011 grams/1 mole
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CALCULATIONS USING THE MOLAR MASS – USE DIMENSIONAL ANALYSIS Calculate the number of grams of
carbon in 3.25 moles of carbon.
Calculate the number of moles of hydrogen in 6.05 grams of hydrogen.
Calculate the number of atoms of carbon in 15.00 grams of carbon.
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MOLE SING A LONG
October 24, 2005
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IT’S A UNIT AFTER ALL”HTTP://WWW.YOUTUBE.COM/WATCH?V=ZNNBZGNSOHK
A mole of laughter, a mole of tears A mole of atoms, a mole of cheer The name of that measure Is a real chemist’s treasure It’s a unit after all Chorus
It’s a unit, after all It’s a unit, after all It’s a unit after all It’s a unit after all
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IT’S A UNIT AFTER ALL” http://www.youtube.com/watch?
v=at_9A_gfln0 A chemist’s friend, tried and ture, An Avogadro would stand by you. And any chemist anywhere, Would stand up and swear, It’s a unit after all Chorus
It’s a unit after all, etc.
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“IT’S A UNIT AFTER ALL” Six point oh two times ten to twenty-
three A number to live by in chemistry So this is October 24th Don’t be absurd, for It’s a unit after all Chorus
It’s a unit after all It’s a unit after all It’s a unit after all I’ts a unit after all