Atoms: The Building Blocks of Matter

PART 1

 Democritus [400 B.C]  Greek philosopher  Hypothesized: Nature has a basic indivisible particle of which everything is made of Called this particle an atom Greek “atomos” = indivisible

 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. H 2 O always has 2 atoms of H and 1 atom of O

 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 CO 2 CO = 1.00g of C and 1.33 g of O CO 2 = 1.00 g of C and 2.66 g of O The ratio of the second element is 2.66 to 1.33 or 2 to 1

 John Dalton [1808]  English schoolteacher – liked nature and weather  Developed: Dalton’s Atomic Theory

1. All matter is composed of extremely small particles called atoms 2. 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

 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

 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 particles  Protons, neutrons, electrons

 J.J. Thomson [1897]  Discovered: The 1 st 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

 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 =7YHwMWcxeX8&NR=1

 Thomson measured the ratio of the charge of the particles to their mass  Same ratio no matter what metal or gas was used  Named this particle an electron 

 Atoms are electrically neutral  Must have positive charges to balance the negatively charged electrons  Electrons have a lot less mass than atoms  Other particles must account for their mass  Plum Pudding Model  positively charged sphere with electrons dispersed through it

 Robert Millikan [1909]  Discovered: The measurement of an electron charge  Oil Drop Experiment Measured the difference in velocity of oil droplets Charged droplets (ionizing radiation) vs. uncharged  /watch?v=XMfYHag7Liw&fe ature=related /watch?v=XMfYHag7Liw&fe ature=related

 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” ww.yout ube.com /watch?v =wzALbz Tdnc8&f eature=r elated

 Discovered the nucleus is a small densely packed volume of positive charge  Size 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

PART 2

 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 = x Neutrons = x To simplify, both have mass of 1 amu (atomic mass unit)

 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!!!

 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 )

 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!!!

 Neutral atoms  total positive charge equals the total negative charge # protons (+1 each) = # electrons (-1 each)

 Atoms of the same element (i.e. same # of protons) that have differing number of neutrons  Isotopes of the same element  have different masses  do not differ significantly in chemical behavior

 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 ElementAtomic Number # of Protons # of Neutrons Mass Number Carbon Nitrogen15

ParticleSymbolChargeMass Number Electrone -, 0 e0 Protonp +, 1 H+11 Neutronn◦, 1 n01

 Atoms with a charge  Negative – more electrons than protons  Positive – more protons than electrons  Charge = #protons - # electrons  Magnesium atom with 12 protons and 10 electrons has a charge of +2

 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 nature  How 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

 There are two naturally occurring isotopes of hydrogen  Hydrogen with 1 proton and zero neutrons  Hydrogen 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 neutrons H amu proton + 1 neutron H amu0.0115

 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 = x of any substance Typically used to talk about particles (atoms, compounds, etc.)

 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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 1 mole = 6.02 x of anything (atoms, molecules, formula units, particles, etc.)  Use dimensional analysis when solving:  Conversion factor: 1 mole = 6.02 x 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 molecules of carbon dioxide, how many moles do I have?

 Molar Mass  The mass of one mole of a substance  The molar mass of an element can be found on the periodic table  Same as the average atomic mass  1 amu = 1 gram/mole E.g. Average atomic mass of C = amus Molar mass of C = grams/1 mole

 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 grams of carbon.

October 24, 2005

 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

 0  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.

 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  I’ts a unit after all