Download Example of calculating average atomic mass

Chem 1 Notes: Chapter 2: The Components of Matter
Categories of Matter: see powerpoint lecture for chart and copy it to your notes.
(So…we identify materials in these four ways: element, compound, homogenous mixture, or
heterogeneous mixture. KNOW THEM!)
Definitions for Components of Matter (from Fig 2.1 parts c and d)
Compound: a substance composed of two or more elements which are chemically combined
Mixture: a group of two or more elements and/or compounds which are physically intermingled
Figure 2.19 The distinction between mixtures and compounds.
Mixture: Physically mixed therefore can be separated by physical means.
Compound: Allowed to react chemically therefore cannot be separated by physical means.
Mixtures
Heterogeneous mixtures: have one or more visible boundaries between the components.
Homogenous mixtures: have no visible boundaries because the componints are mixed as
indivdial atoms, ion, and molcules.
Solutions: a homogenous mixture is also called a solution. Solutions in water are called
aqueous solutions and are very important in chemistry.
Basic Separation Techniques: filtration, crystallization, distillation, extraction, chromatography
Figure B2.3 & B2.4 from 4th ed: Filtration and Crystallization
PROPERTIES OF MATTER
Physical Properties: color, texture, odor, density, BP, MP, etc.
Chemical Properties: how a substance reacts with other substances
Table 2.1 Some Properties of Sodium, Chlorine and Sodium Chloride: chemical property is the
reaction of sodium with chlorine to make sodium chloride (and also when Na reacts with water rather
than just dissolving in it); physical properties of Na, Cl and NaCl listed: MP, BP, color, density,
solubility of Cl2 and NaCl in water.
Be able to:
Identify the type of matter in four ways: element, compound, homogenous mixture, or
heterogeneous mixture
Know the three states of matter
Separate physical properties from chemical properties
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Law of conservation of mass: the total mass of substances does not change during a chemical
reaction. (In a chemical reaction, the mass of reactants used has to equal the mass of products formed.)
reactant 1 + reactant 2 + …  product(s)
total mass in
= total mass out
CaO + CO2  CaCO3
56.08g + 44.00g = 100.08 g
Law of Definite (or Constant) Composition: No matter the source, a particular compound is
composed of the same elements in the same parts (fractions) by mass. (See table below for calcium
carbonate.)
Analysis by mass (g/20.0g)
8.0 g Ca
2.4 g C
9.6 g O
20.0 g
Mass fraction
0.40 Ca
0.12 C
0.48 O
1.00
Mass Percent
40. %-mass Ca
12 %-mass C
48 %-mass O
100. %-mass
Law of Multiple Proportions: If elements A and B react to form two compounds, the different
masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers.
Example: There are two “Carbon Oxides,” A & B
Carbon Oxide I : 57.1% oxygen and 42.9% carbon
Carbon Oxide II : 72.7% oxygen and 27.3% carbon
Assume that you have 100 g of each compound. That means that Oxide I has 57.1 g of oxygen and 42.
9 g of carbon. Oxide II has 72.7 g of oxygen and 27.3 g of carbon.
Mass ratio of oxygen to carbon for Oxide I: 57.1 g oxygen/42.9 g carbon = 1.33
Mass ratio of oxygen to carbon for Oxide II: 72.7 g oxygen/27.3 g carbon = 2.66
2.66 for Oxide II/1.33 for Oxide I = 2/1 ratio; i.e., twice as much oxygen in Oxide II
Dalton’s Atomic Theory (Memorize!)
1. All matter is composed of atoms.
2. Atoms of one element cannot be converted into atoms of another element.
3. Atoms of an element are identical in mass and other properties and are different from atoms of any
other element.
4. Compounds result from the chemical combination of a specific ratio of atoms of different elements.
Fig 2.4 Experiments to determine the properties of cathode rays. (Conclusions: cathode rays consist of
charged particles, they are negative, and the particles are found in all matter.)
Fig 2.5 Millikan’s oil-drop experiment for measuring an electron’s charge.
Millikan used his finding to also calculate the mass of an electron. (9.109x10-28g)
Early Atomic Theory
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At the beginning of the 20th Century: scientists knew an atom contained both positive and
negative particles. One model was “plum pudding” – like raisins in rice pudding.
Rutherford’s experiment showed that an alpha particle (a Helium nucleus with 2 protons and 2
neutrons) aimed at gold foil was deflected almost straight back. NOT WHAT HE EXPECTED: HAD
TO REVISE THEIR MODEL!!! (See Figure 2.6)
Led to modern nuclear atom model.
Figure 2.6 Rutherford’s Experiment: A. Hypothesis: Expected result based on “plum pudding” model
and
minor deflections are seen. C. Actual Result: major deflection indicates atoms have a tiny, massive,
positive nucleus.
Figure 2.7 General features of a nuclear atom (Know these: atom is electrically neutral, spherical &
composed of a positively charged central nucleus surrounded by one or more negatively charged
electrons. Nucleus consists of protons and neutrons.)
Table 2.2 Properties of the Three Key Subatomic Particles (You need to MEMORIZE names,
symbols, relative charge, relative atomic mass in amu to sig figs shown in this table, and location in
atom)
Atomic Symbols, Isotopes, Numbers
Z
AX
X is the atomic symbol of the element; A is mass number which is sum of Z + N; Z is atomic number
or the number of protons in the nucleus; N is the number of neutrons.
Isotope: atoms of an element with the same number of protons but a different number of neutrons.
Sample Problem 2.4 and you also try these uranium isotopes 234, 235, 238, and 239 (answer here).
ATOMS & ISOTOPES
REMEMBER: In an atom of the same element:
# protons never changes (except for nuclear decay)
# neutrons can be different ==> ISOTOPES
# electrons can change ==> IONS
Figure 2.9 The Mass Spectrometer and its Data (Actual mass and percent abundance are found for each
isotope of an atom.)
AVERAGE ATOMIC MASS
By using a mass spectrometer, an atom’s isotopes can be counted, establishing relative abundance of
each isotope. This is then used to calculate a weighted average atomic mass:
Weighted average atomic mass =
fract1*mass1 + fract2*mass2 + ...
Example of calculating average atomic mass
We put boron in a mass spectrometer, and find there are just two isotopes, with these results: 19.91%
Boron-10 mass 10.0129 amu and 80.09% Boron-11 mass 11.0093 amu.
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Average atomic mass = 0.1991 * 10.0129 + .8009 * 11.0093
= 1.994 + 8.817 = 10.811 (sig fig rules)
Do problem 33 for practice on this method. (Magnesium has three naturally occurring isotopes. 24Mg
has a mass of 23.9850 amu and 78.99% abundance, 25Mg has a mass of 24.9868 amu and 10.00%
abundance, and 26Mg has a mass of 25.9826 amu and 11.01% abundance.)
Difference between atomic mass and mass number
You have just practiced determining the weighted average atomic mass. How is that different from
mass number?
ATOMIC MASS IS NOT THE SAME AS MASS NUMBER!
ATOMIC MASS IS NOT THE SAME AS MASS NUMBER!
ATOMIC MASS IS NOT THE SAME AS MASS NUMBER!
Think you can remember this?
The Modern Reassessment of the Atomic Theory:
1. All matter is composed of atoms. The atom is the smallest body that retains the unique identity of
the element.
2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction.
Elements can only be converted into other elements in nuclear reactions.
3. All atoms of an element have the same number of protons and electrons, which determines the
chemical behavior of the element. Isotopes of an element differ in the number of neutrons, and thus in
mass number. A sample of the element is treated as though its atoms have an average mass.
4. Compounds are formed by the chemical combination of two or more elements in specific ratios.
Figure 2.10 The modern periodic table. Know that the Main Group (or Representative) Elements are
columns 1-2 and 13-18, the Transition Elements are columns 3-12, and the Inner Transition Elements
fit inside the Transition Elements. Also know the location of the metals, nonmetals and metalloids.
You need to memorize the first 36 elements’ symbol and names, plus ten more: Ag, Sn, I, Ba, Pt, Au,
Hg, Pb. Rn, U.
Tell me what you know about groups, columns, trends, etc.
Write down group names, locations for metal, nonmetals and metalloids.
What trends to you know?
Where are the transition metals and the inner transition metals?
Figure 2.11 (4th ed.) Metals, metalloids, and nonmetals
****Go to Chapter Seven***************************************
Section 2.7 goes with chapter 9:
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Figure 2.14 (4th ed.) The relationship between ions formed and the nearest noble gas
Figure 2.13 Formation of a covalent bond between two H atoms
****Section 2.8 goes after chp 11
Use the lecture notes, the textbook AND the lab manual’s Dry Lab on Nomenclature.
All the tables in the dry lab will be very useful!
You must memorize these!
The seven common diatomic elements: H2, O2, N2, F2, Cl2, Br2, I2 (HON + Halogens)
Types of Chemical Formulas
A chemical formula is comprised of element symbols and numerical subscripts that show the type and
number of each atom present in the smallest unit of the substance.
Empirical formula indicates the relative number of atoms of each element in the same compound. It
is the simplest type of formula. The empirical formula of hydrogen peroxide is HO.
Molecular formula show the actual number of atoms of each element in a molecule of a compound.
The molecular formula for hydrogen peroxide is H2O2.
A structural formula show the number of atoms and the bonds between them. H-O-O-H
Figure 2.16 Some common monatomic ions of the elements
Naming metal ions with more than one oxidation number (charge)
Remember that metals like iron or tin with more than one possible charge have to show which one by
using Roman numerals for the charge in parentheses.
Sn2+ is tin (II) ion
Sn4+ is tin (IV) ion
Naming binary ionic compounds: see rules in your packet.
The name of the cation is written first, followed by that of the anion.
The name of the cation is the same as the name of the metal. Many metal names end in –ium
The name of the anion takes the root of the nonmetal name and adds the suffix –ide.
Calcium ion and bromide ion form calcium bromide.
Example: form an ionic compound of Al & Br and then Al & O.
Al forms Al3+ and Br forms Br-, so the ration will be 1:3 or AlBr3, aluminum bromide
Again, Al3+ and O forms O2-, so the ratio will be 2:3 or Al2O3, aluminum oxide
Special trick called the criss-cross rule. Xa+ + Yb-  XbYa, but if b=a, reduce formula to 1:1 ratio
Sample Problem 2.7 & 2.8 combined plus follow-up problems.
Sample Problem 2.9
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Hydrates of ionic compounds
Hydrates are ionic crystals of salts with water molecules incorporated in their crystal structures.
Write “formula unit name – dash – Greek prefix (representing # of water molecules) hydrate”
BaCl2.2H2O is barium chloride dihydrate.
You try:
Name CuSO4.5H2O:
Give the formula for sodium sulfate decahydrate:
Figure 2.17 Naming Oxoanions and Table 2.6 Numerical Prefixes
Sample Problem 2.10 Determining Names and Formulas of Ionic Cmpds Containing Polyatomic Ions
Naming Inorganic Acids – see rules in Dry Lab II in lab manual.
1) Binary acid solutions form when certain gaseous compounds dissolve in water.
For example, when gaseous hydrogen chloride (HCl) dissolves in water, it forms a solution
called hydrochloric acid. Prefix hydro- + anion nonmetal root + suffix -ic + the word acid hydrochloric acid
Practice naming H2S(aq):
2) Oxoacid names are similar to those of the oxoanions, except for two suffix changes:
Anion “-ate” suffix becomes an “-ic” suffix in the acid. Anion “-ite”
suffix becomes an “-ous” suffix in the acid.
The oxoanion prefixes “hypo-” and “per-” are retained. Thus, BrO4is perbromate, and HBrO4 is perbromic acid; IO2- is iodite, and
HIO2 is iodous acid.
Practice naming HClO3.
Sample Problem 2.12: name the acids formed when these anions gain enough H+ to become neutral.
a. Brb. IO3c. CNd. SO42e. NO2Naming Binary Covalent Compounds
Very simple – you will know a compound is covalent if it’s two nonmetals. Indicate how many
of each atom in the compounds by using Greek prefix.
If first element is single, leave off mono. If first element is hydrogen, leave off any prefix.
‘Prefix’element name – ‘prefix’element root – suffix ‘ide’
Practice: CO, CO2, NO2, N2O4, P2O5, HF, H2S
Common names to memorize: water H2O, hydrogen peroxide H2O2, ammonia NH3, hydrazine N2H4,
NO nitric oxide, N2O nitrous oxide.
Organic Formulas and Names
Memorize: methane CH4, ethane CH3CH3, propane CH3CH2CH3, butane CH3(CH2)2CH3,
octane CH3(CH2)6CH3, benzene C6H6, methanol CH3OH, ethanol CH3CH2OH, 1-propanol
CH3CH2CH2OH, acetic acid CH3COOH, formaldehyde HCHO, glucose C6H12O6, sucrose C12H22O11.
This part goes with or just before chapter 3:
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Calculating molar mass of compounds: look up the molar mass on the Periodic Table. Sum up all
atoms in the compound.
Sample Problem 2.15
PRACTICE:
Follow-up Problem 2.15: Determine the formula and the molecular (or formula) mass for each of
these:
a. hydrogen peroxide
b. cesium chloride
c. sulfuric acid
d. potassium sulfate
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