Atoms and Elements Chapter 2 1

Atomic theory  John Dalton, 1808: matter is made of tiny indestructible particles called atoms  What evidence persuaded John Dalton that matter was made of atoms?

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Evidence for Atoms  Boyle’s Law (1660s) • • Gases can be compressed (PV = constant) Suggests that a gas is made of particles with space between  Conservation of mass (Lavoisier, 1789) • In a chemical reaction, matter is neither created nor destroyed 3

Evidence for Atoms  Law of Definite Proportions (Proust, 1797) • All samples of a compound, regardless of source or how prepared, have the same proportions of their constituent elements (constant composition)  Law of Multiple Proportions (Dalton, 1803) • When two elements (A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small whole numbers.

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Law of Multiple Proportions 2.67

gO gC

1.33

gO gC

 2 1 5

Evidence for Atoms  Combining Volumes of Gases • When two gases combine to form a new compound, the volumes that combine will be in a ratio of small whole numbers.

hydrogen nitrogen

 3 1

hydrogen



oxygen

2 1

hydrogen chlorine

 1 1 6

A New System of Chemical Philosophy: John Dalton, 1808  Each element is composed of tiny indestructible particles called atoms.

 All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements.

 Atoms combine in simple, whole-number ratios to form compounds.

 Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms change the way that they are bound together with other atoms to form a new substance.

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Cathode ray tube 8

Discovery of the electron  J.J. Thompson (1897) • Cathode rays are deflected by an electric or magnetic field • Cathode ray = beam of negatively charged particles (electrons) coming out of cathode metal atoms • Determined electron mass-to-charge ratio = –5.6857 x 10 –9 g/C 9

Millikan oil-drop experiment 10

Charge of the electron  Robert Millikan (1909)  Drop charges are integer multiples of 1.60 x 10 –19 Coulomb 11

Radioactivity  Some elements spontaneously emit radiation called radioactivity  Three types of radioactivity: • Alpha ( a ) = heavy, +2 charge • Beta ( b ) = light, –1 charge • Gamma (g = high-energy photon 12

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The gold foil experiment  Ernest Rutherford (1910)  Bombarded gold foil with alpha particles • Most went straight through • Some slightly deflected  • A few strongly deflected “about as credible as if you had fired a 15-inch shell at a piece of tissue paper, and it came back and hit you” 14

Alpha particle scattering explained 15

Rutherford’s nuclear model 10 -10 m 10 -15 m  Mass & positive charge concentrated in nucleus  Most a particles miss the nucleus & are not deflected:  most of the atom is empty space Some a particles come near the nucleus & are deflected: nucleus is positively charged and very dense  Tiny, lightweight electrons circle the nucleus, like planets around the sun 16

Composition of the atom  Protons • • +1 charge 1.67262 x 10 –24 g = 1.0073 amu  Neutrons • • 0 charge 1.67493 x 10 –24 g = 1.0087 amu  Electrons • • –1 charge 0.00091 x 10 –24 g = 0.00055 amu 17

Atomic number and mass number

 Atomic number ( Z ) = protons  Mass number ( A ) = protons + neutrons  A 35 Cl atom has 17 protons and 18 neutrons  Also shown as chlorine-35 18

Ion charge

 When an atom loses or gains an electron, it becomes an ion  Ion charge is shown in the upper right corner of the atomic symbol  If no charge is shown, the charge is zero mass number, A protons + neutrons atomic number, Z protons 23 11 Na 1+ ion charge protons – electrons 19

Give the structure (p + , n 0 , e – ) of these atoms 19 F 40 K 1+ 235 U 34 S 2–  9 p + , 10 n 0 , 9 e –  19 p + , 21 n 0 , 18 e –  92 p + , 143 n 0 , 92 e –  16 p + , 18 n 0 , 18 e – 20

Give the symbol of each atom or ion  35 p + , 44 n 0 , 36 e –  79 Br 1–  47 p + , 62 n 0 , 46 e –  109 Ag 1+  26 p + , 28 n 0 , 26 e –  54 Fe 21

Isotopes

 Atoms with the same number of protons but different numbers of neutrons are isotopes of the same element  6 Li and 7 Li are isotopes of lithium    Both are the element lithium 6 Li has 3 protons, 3 neutrons 7 Li has 3 protons, 4 neutrons 22

A Mass Spectrometer 23

02_12 Fractional abundance Mass Spectrum of Neon 20 Ne 90.48% 21 Ne 22 Ne 0.27% 9.26% 18 19 20 21 Atomic mass (amu) 22 24

Average atomic mass  Element’s atomic mass = weighted average of masses of all naturally-occurring isotopes of that element 25

Average Atomic Mass  Calculate the average atomic mass of Ne:  ISOTOPE ISOTOPIC MASS (amu) ABUNDANCE 20 Ne 21 Ne 22 Ne 19.99244 amu 20.99395 amu 21.99138 amu 90.48% 0.27% 9.26% 26

Average Atomic Mass  Calculate the average atomic mass of Ne:  ISOTOPE ISOTOPIC MASS (amu) ABUNDANCE 20 Ne 21 Ne 22 Ne 19.99244 amu 20.99395 amu 21.99138 amu 90.48% 0.27% 9.26% (0.9048)(19.99244 amu) = 18.09 amu (0.0027)(20.99395 amu) = 0.057 amu (0.0926)(21.99138 amu) = 2.04 amu 20.18 amu 27

Atomic mass  Two natural isotopes of antimony exist. 57.3% exists as 121 Sb (mass 120.9038 amu), and the rest is 123 Sb (mass 122.9041 amu). What is the atomic mass of antimony?

 The abundances must total 100% , so abundance of 123 Sb = 100 – 57.3 = 42.7%  (0.573 x 120.9038) + (0.427 x 122.9041) = 69.3 + 52.5 = 121.8 amu 29

Percent abundance  Two natural isotopes of copper exist, 63 Cu (62.9296 amu) and 65 Cu (64.9278 amu). What is the abundance of each isotope?

 The abundances must total 100%  If x = 63 Cu abundance, then 65 Cu = 1–x  Average mass (from periodic table) = 63.546 amu 30

Percent abundance  x = 63 Cu (62.9296 amu) & 1–x = 65 Cu (64.9278 amu)  Average mass = 63.546 amu  (  ( 62.9296

amu

    ( 1 

x

 ( 64.9278

amu

   63.546

amu

62.9296

x

 64.9278

 64.9278

x

 63.546

 1.9982

x

  1.382

x

  1.382

 1.9982

 0.6915 1 

x

 0.3085

 63 Cu = 69.15% and 65 Cu = 30.85% 31

A unit for counting atoms  We have units for describing mass, volume, temperature, and so forth, but  we need a unit for counting the number of items  For large items like eggs or donuts, we can use dozen: one dozen = 12 items  For tiny items like atoms or molecules, we need . . . the mole!

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What is a mole?

 One mole is defined as the number of atoms in exactly 12 grams of 12 C  One mole contains 6.0221421 x 10 23 particles  This number is called Avogadro’s number (N A )

There is Avogadro’s number of particles in a mole of

any substance

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Calculations with Avogadro  How many atoms of gold are present in 0.0507 mol Au?

0.0507

mol

 6.022

 10 23

atoms

1

mol

 3.05

 10 22

atoms

 How many moles of Pb atoms are 8.27 x 10 22 atoms of Pb? 8.27

 10 22

atoms

 6.022

1

mol

 10 23

atoms

 0.137

mol

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Mass on the periodic table One atom of Cl weighs 35.45 amu OR One mole of Cl atoms weighs 35.45 grams The molar mass of Cl is 35.45 g/mol 35

Calculations with molar mass  What is the mass of 1.38 mol Al?

1.38

mol

 26.98

g

1

mol

 37.2

g

 How many moles are in 35 g of Zn?

35

g

 1

mol

65.39

g

 0.54

mol

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Avogadro and molar mass  What is the mass of 2.35 x 10 24 atoms of Cu?

2.35

 10 24

atoms

 6.02

1

mol

 10 23

atoms

 3.90

mol Cu

3.90

mol Cu

 63.55

g

1

mol

 248

g Cu

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Avogadro and molar mass  How many He atoms are present in a 22.6 g sample of He gas?

22.6

g He

 1

mol

4.0026

g

 5.65

mol He

5.65

mol

 6.02

 10 23

atoms

1

mol

 3.40

 10 24

atoms

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