Chapter 2.1-2.4 & definition of moles

Chapter 2.1-2.4 & definition of molesAtoms, Molecules, and Ions

Atomic Theory (Dalton)

1. Elements are made up of atoms (Na, K, H, etc.)

2. Atoms of same elements have same properties

3. Atoms are neither created nor destroyed in a chemical reaction (Law of Conservation of mass)

4. Compounds are formed by combinations of elements in fixed ratios.

H2O 2 H, 1 O

• What makes an atom?– How can one get that information?

• What makes an egg?– How can you get that information?

ATOM

Electron (e) Proton (p) Neutron (n)

Charge -1 +1 0(au)

Mass 1/1837 1 1(au)

1 amu = 1.66 x 10-24 g

Discovery & Characterization of the

electron (e)

• Cathode Ray (-) charges electrons

Thompson’s Experiment

• With electrically charged plates and a magnet

• Charge/mass ratio for an electron– (e/m) = 1.76 x 108 Coulomb/g (or

C/g)

Millikan oil-drop experiment

Charge of e- = 1.60 x 10-19 Coulomb (C)Mass of e- = 9.10 x 10-28 g

Radioactivity• Spontaneous emission of

radiation

α ( )

β ( ) electron

γ ( γ)

• Note: 238 is # protons + # neutrons, 92 is # protons

• α is attracted to (-) plate and hence must contain (+) charge

23892U

42He

01e

00

23892U

Rutherford’s Experiment

• Atom– Protons (p), Neutrons (n), and

Electrons (e-)•How are they together?•Are they uniformly spread over?

Rutherford’s Experiment

•α-Particles are allowed to strike a gold plate

»Most particles go through»A few go backward

Rutherford’s Experiment

• Can we predict what is going on in Rutherford’s Experiment?

• Why do α –particles (used as balls) go through undeflected?

Rutherford’s Experiment

• Conclusion– An atom is mostly made up of

empty space (like an open door) with a massive nucleus (like a concrete wall), occupying a very small volume.

– Nucleus has protons and neutrons, electrons revolve outside.

Isotopes, Atomic Number, and Mass Number

• Atomic number = Number of Protons

• Mass number = # Protons + #

Neutrons

• Can you define “isotopes”?– Same elements?– How are they different?

Isotopes, Atomic Numbers, and Mass Numbers

•Here is an example of Isotopes. •Remember that the top number is the mass # and the bottom number is the atomic #.

11H

21H 3

1H

Average Atomic Mass• Avg. mass = P1M1 + P2M2 + …

– P1 = % abundance for Isotope 1

– M1 = Mass for Isotope 1

– P2 = % abundance for Isotope 2

– M2 = Mass for Isotope 2

Refer to Problem 29 (Pg. 72)

Mass Spectrometer• Obtain atomic and molecular

weights • Need gaseous samples• Ionization of sample by high

energy electrons• Pass gaseous ions through

poles of a magnet• For ions carrying the same

charge, more massive particles will be less deflected than the less massive particles. Thus, mass separation takes place.

Mole & Number• 1 thousand = 103 = 1000• 1 million = 106

• 1 mole = 6.022 x 1023 (Avogadro’s number)

• 1 mole H atoms = 6.022 x 1023 H atoms

• 1 mole H2O molecules = 6.022 x 1023 H2O molecules

Mole & Gram (Molar Mass)

(Remember: M-g-P = mole-gram-Periodic Table)

• 1 mole C weighs 12.0107 g ≈ 12.0 g C atoms – Found on the Periodic Table

• 1 mole H weighs 1.00794 g ≈ 1.00 g H atoms

• 1 mole H2O = 2(1.00794) + 15.9994 ≈ 2 + 16 = 18 g H2O molecules