Download Which has more atoms: a one gram sample of carbon

Chem 152 Chapter 5 Part 1: Early Atomic Theory and Structure.
Sections 5.1 – 5.5
Dalton’s Atomic Theory
 Each element consists of indivisible*, minute particles called atoms.
 Atoms can neither be created or destroyed in chemical reactions.
 All atoms of a given element are identical*.
 Atoms combine in whole-number ratios to form compounds.
 Atoms of different elements have different masses.
 Numbers 1 and 3 are not completely correct!
What’s Inside the Atom?
 Atoms are not indivisible!
 There are three types of “subatomic” particles inside each atom:
 Proton
 Neutron
 Electron
 Determined by many scientists, experiments over many years.
Protons, Neutrons and Electrons
 Protons and neutrons are much heavier than electrons (1800x heavier).
 Electrical charge:
 Protons “positive” +
 Electrons “negative” –
 Neutrons “neutral”
 Protons and neutrons are located in the center of the atom in a densely packed
“nucleus”.
 Electrons “fly around” outside the nucleus in a cloud-like area.
Results of Rutherford’s Experiment
 Most of the mass of the atom is confined to a dense nucleus.
 The nucleus has a positive charge.
 The rest of the atom is mostly empty space.
Protons Define an Element
 All atoms of the same element have the same number of protons.
 Hydrogen: 1 proton
 Carbon: 6 protons
 Oxygen: 8 protons
 Iron: 26 protons
 Lead: 82 protons
 The atomic number of an element is the same as its number of protons.
Isotopes
 Atoms of the same element may have different numbers of neutrons.
 Carbon may have 6, 7 or 8 neutrons.
 Hydrogen may have 0, 1 or 2 neutrons.
 These are called isotopes.
 Most elements have more than one isotope.
 Some isotopes are radioactive.
 Unstable, decay into other elements.
Example: Isotopes of Carbon
 Remember:
 Carbon (C) has an atomic number of 6.
 All C atoms have 6 protons.
 The proton and neutron both have a relative mass of 1.
 C may have 6, 7 or 8 neutrons.
 Therefore, the relative mass of a C atom will be 12, 13 or 14 depending on how
many neutrons it has.
 This is called the mass number.
 These isotopes of carbon are called carbon-12, carbon-13 and carbon-14.
 We label them 12C, 13C and 14C.
Which has more atoms: a one gram sample of carbon-12, or a one gram sample
of carbon-13?
Isotopes of Hydrogen
 hydrogen-1: protium
 hydrogen-2: deuterium
 hydrogen-3: tritium
Mass Number
 To get the mass number of an isotope, simply add up the number of protons and
neutrons.
 An atom with 11 protons and 12 neutrons has what mass number?
 What is the element?
Complete the Following:
Given the isotope 67Zn:
Atomic # ?, Mass # ?, # of protons, # of neutrons, # of electrons
Atomic Mass
 The total mass of an atom’s components (protons, neutrons and electrons) is
call its atomic mass.
 Units: atomic mass unit: amu
 Each isotope of an element has a particular “natural abundance”.
 The percentage of each isotope found in a natural sample of the element.
12

C has a natural abundance of 98.89%.
13
 Most of the rest is
C.
 Atomic masses listed in the periodic table are averages of the atomic masses of
each isotope.
Atomic Mass Units
 1/12 the mass of a carbon-12 atom.
 1 amu = 1.6606 x 10 –24 g.
 All elements have atomic masses that are relative to carbon-12.
Calculation of Average Atomic Mass (weighted average mass)
 Calculate the atomic mass of copper given the following data:
 63Cu: 62.930 amu, 69.09% abundance
 65Cu: 64.928 amu, 30.91% abundance
 (over all isotopes) fraction abundance x (exact weight)
average mass = (0.6909)(62.930) + (0.3091)(64.928)
Homework
 p. 131, Key Concepts: 2, 3, 4
 p. 132, Exercises: 1 – 4, 11, 13, 15 – 27, 33 – 37.