Chapter 2 - Atoms, Molecules, & Ions Flashcards
-Describe Dalton's atomic theory -Describe the key experiments that led to the discovery of electrons & to the nuclear model of the atom -Describe the structure of the atom (proton, neutron, electron) -Describe the electric charge & relative masses of atomic particles -Use chemical symbols w/ atomic # & mass # to express the subatomic composition of isotopes -Understand atomic weights & how they relate to the masses of individual atoms & their natural abundances -Describe how elements are
2.1
Define “Dalton’s atomic theory”.
- Each element is composed of atoms.
- All atoms of a given element are identical in mass & other properties, but atoms of one element are different from the atoms of others.
- The atoms of one element cannot be changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
- Compounds are formed when atoms of >1 element combine; a given compound always has the same relative # & kind of atoms.
2.1 - Give It Some Thought (p. 39)
1 compound of C & O contains 1.333g of O per gram of C, whereas a second compound contains 2.666g of O per gram of C.
(a) What chemical law do these data illustrate?
(b) If the 1st compound has an equal # of O & C atoms, what can we conclude about the composition of the 2nd compound.
(a) Law of multiple proportions
(b) The 2nd compound must contain 2 O atoms for each C atom
2.1
Define “Law of Constant Composition”.
In a given compound, the relative numbers and kinds of atoms are constant.
Basis of Dalton’s Postulate 4.
2.1
Define “Law of Conservation of Mass”.
The total mass of materials present after a chemical reaction is the same as the total mass present before the reaction.
Basis of Dalton’s Postulate 3.
–Atoms always retain their identities & that atoms taking part in a chemical reaction rearrange to give new chemical combinations.
2.1
Define “Law of Multiple Proportions”.
If 2 elements A + B combine to form >1 compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers.
There are compounds that use the same elements, but are not the same compound due to the amount.
Ex. H2O & H2O2
- H2O: 1.0g H + 8.0g O
- H2O2: 1.0g H + 16.0g O
Ex. CO & CO2 … small whole numbers
2.2
J.J. Thomson - 1897
- Experiment
- Discovery
- Model
- Cathode Ray experiment.
- Discovered electrons.
- Plum pudding model - + charged sphere with electrons scattered throughout
2.2
Robert Millikan - 1909
- Experiment
- Discovery
Measured the charge of an electron.
Small drops of oil, which had picked up extra electrons, were allowed to fall btwn. 2 electrically charged plates.
1. Millikan monitored the drops, measuring how the voltage on the plates affected their rate of fall.
- Charges were always integral multiples of 1.602 x 10^-19 Coulombs - the charge of a single electron.
Discovered mass by dividing charge by charge-to-mass ratio. Mass is 9.10 x 10^-28g.
2.2
Henri Becquerel - 1896
Discovered radioactivity while studying a U compound
2.2
Ernest Rutherford
1.Gold foil experiment. Aimed alpha particles at gold tin foil. Most went straight through, but some deflected. Postulated that mots of the mass of each Au atom in his foil & all of its + charge reside in a very small, extremely dense region- nucleus.
- Nucleus, Protons
Alpha, beta, gamma radiation
2.2
James Chadwick - 1932
Neutrons
2.3
Why do atoms lack a net electrical charge?
Every atom has an equal number of electrons & protons.
2.3 Give It Some Thought (p. 43)
(a) If atom has 15 protons, how many electrons does it have?
(b) Where do the protons reside in an atom?
(a) 15
(b) Nucleus
2.3 Sample Ex.
The diameter of a US penny is 19 mm. The diameter of an Ag atom is only 2.88 A. How many Ag atoms could be arranged side by side in a straight line across the diameter of a penny?
(19 mm)(10^-3 m / 1 mm)(1 A / 10^-10 m)(1 Ag atom / 2.88 A) = 6.6 x 10^7 Ag atp,s/
2.3 Practice Ex.
The diameter of a C atom is 1.54 A.
(a) Express this diameter in pm.
(b) How many C atoms could be aligned side by side in a straight line across the width of a pencil line that is .20 mm wide?
a. (1.54 A)(10^-10 m / 1 A)(10^12 pm / 1 m) = 154 pm
b. (.20 mm)(10^-3 m / 1 mm)(1 A / 10^-10 m)(1 C atom / 1.54 A) = 1.3 x 10^12 C atoms
In calc… .2xE^-3/E-10/1.54
2,3
Define “atomic number”.
The number of protons in the nucleus of an atom of any particular element.
2.3
Define “mass number”.
The total number of protons + neutrons in the atom.
2.3
Indicate 3 ways to write Carbon-12.
12C ; carbon-12; 12C
6
2.3
Define “isotope”.
Atoms with identical atomic numbers, but different mass numbers (same proton #, different neutron #).
2.3 - Samp. Ex.
How many P, N, & E are in
(a) 1 atom of 197 Au
(b) 1 atom of Sr-90
(a) 79 P, 79 E, 118 N
(b) 38 P, 38 E, 52 N
2.3 - Samp. Ex.
Mg has 3 isotopes, w/ mass numbers 24, 25, and 26.
(a) Write the complete chemical symbol for each of them.
(b) How many neutrons are in an atom of each isotope?
(a)
24Mg ; 25Mg ; 26Mg
12 12 12
(b)
12, 13, 24
2.4
1 amu = x g
1.66054 x 10^-24g
2.4
1g = x amu
6.02214 x 10^23 amu
- 4
- 93% C-12; 1.07% C-13
Find amu.
(.9893)(12) + (.0107)(13) = 12.01 amu
2.4
Average atomic mass is also known as
Atomic weight
2.4 - Give It Some Thought (p. 47)
A particular atom of Cr has a mass of 52.94 amu, whereas the atomic weight of Cr is 51.99 amu. Explain the difference in the two masses.
Cr has different isotopes that weigh less than 52.94, and when you average them all together they come to 51.99.
2.4 Samp. Ex.
Naturally occurring Cl is 75.78% Cl-35, which has an atomic mass of 34.969 amu, & 24.22% Cl, which has an atomic mass of 36.966 amu. Calculate the average atomic mass of Cl.
(.7578)(34.969) + (.2422)(36.966) = 35.45 amu
2.4 Practice Ex.
Three isotopes of Si occur in nature: Si-28 (92.23%) which has an atomic mass of 27.97693 amu; Si-29 (4.68%) which has an atomic mass of 28.97649 amu; and Si-30 (3.09%) which has an atomic mass of 29.97377 amu. Calculate the atomic weight of Si.
(.9223)(27.97693) + (.0468)(28.97649) + (.0309)(29.97377) = 28.09 amu
2.5
What is the name of the metals in Group 1A?
Alkali Metals
Ex. Lithiujm
2.5
What is the name of the metals in Group 2A?
Alkaline Earth Metals
Ex. Calcium
2.5
What is the name of the nonmetals in Group 6A?
Chalcogens
Ex. Oxygen
2.5
What is the name of the nonmetals in Group 7A?
Halogens
2.5
What is the name of the nonmetals in Group 8A?
Noble gases (rare gases)
2.5
What are the metalloids?
B, Si, Ge, As, Sb, Te
2.5
Is H a metal?
No, it’s a nonmetal.
2.5
List metal characteristics.
- Luster
- High electrical & heat conductivity
- Solids at room temperature (except Hg)
2.5 - Give It Some Thought (p. 51)
Chlorine is a halogen. Locate this element in the periodic table.
(a) What’s its symbol?
(b) In what period & group is the element located?
(c) What is its atomic number?
(d) Is it a metal or nonmetal?
(a) Cl
(b) 2, 7A
(c) 17
(d) Nonmetal
2.5 Samp. Ex.
Which two of the following elements would you expect to show the greatest similarity in chemical & physical properties: B, Ca, F, He, Mg, P?
Mg & Ca. Both Alkaline Earth Metals.
2.5 Practice Ex.
Locate Na & Br on the periodic table. Give the atomic number of each, and label metal, metalloid, or nonmetal.
Na: 11, Metal
Br: 35, Nonmetal
2.6
Define “molecule”.
An assembly of 2+ atoms tightly bound together.
2.6
What is the chemical formula of oxygen?
O2
2.6
Define “diatomic molecule.” What elements tend to occur as diatomic molecules?
A molecule that’s made up of 2+ atoms.
H2, N2, O2, and the Halogens (F2, Cl2, Br2, I2)
2.6
Define “molecular compound”.
Compound that consists of molecules.
Ex. H2O.
Ex. H2O2.
2.6
Most molecular substances contain only _____.
Nonmetals
2.6
Define “molecular formula.”
Chemical formula that indicates the actual numbers and types of atoms in a molecule.
Ex. H2O2
Ex. C2H4
2.6
Define “empirical formula.”
Chemical formula that gives only the relative number of atoms of each type in a molecule.
Ex. HO
Ex. CH2
2.7 Samp. Ex. 2.7
Give the chemical symbol, including mass number, for each of the following ions:
(a) The ion w/ 22 P, 26 N, 19 E
(b) The ion of S that has 16 N & 18 E
(a) 48 Ti +3
22
(b) 32 S -2
16
2.7 Prac. Ex.
How many P, N, & E does 79 Se -2 have?
P: 34
E: 36
N: 45
2.7 Samp. Ex.
Predict the charge expected for the most stable ion of Ba & O.
Ba: 2+.
It’s closest noble gas is Xe, which is 2 spots away to the left. Ba gains the Xe configuration when it loses 2 electrons.
O: -2
It’s closest noble gas is Ne, which is 2 spots away to the right. O gains Ne’s configuration when it gains 2 electrons.
2.7 Prac. Ex.
Predict the charge expected for the most stable ion of Al & F.
+3, -1
2.7
Define “ionic compound”.
Compound that contains both positively & negatively charged ions.
2.7
Define “ion”.
A charged particle
2.7
Define “polyatomic ion”.
Consist of atoms joined as in a molecule, but have a net positive or negative charge.
Ex. NH4+ , SO4 -2
2.7
_____ compounds are generally combinations of metals & nonmetals, while _____ compounds are generally composed of nonmetals only, as in H2O.
Ionic, molecular
2.7 Samp. Ex
Which of the following compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4
Na2O and CaCl2. They have metals combined with nonmetals.
2.7 Give It Some Thought (p. 58)
Why don’t we write the formula for the compound formed by Ca 2+ & O 2- as Ca2O2?
Because this is a molecular formula, not an empirical formula. The empirical formula is CaO.
2.7 Samp. Ex.
What are the empirical formulas of the compounds formed by
(a) Al 3+ & Cl- ions
(b) Al 3+ & O 2- ions
(c) Mg 2+ & NO3- ions
(a) AlCl3
(b) Al2O3
(c) Mg(NO3)2
2.7 Prac. Ex.
Write the empirical formulas for the compounds formed by the following ions:
(a) Na + & PO4 -3
(b) Zn 2+ & SO4 2-
(c) Fe 3+ & CO3 2-
(a) Na3PO4
(b) ZnSO4
(c) Fe2(CO3)3
2.8
If a metal can form different cations, the positive charge is indicated by a ______ following the name of the metal.
Roman numeral in parentheses
Ex. Fe 2+ … Iron (II) Ion
Ex. Cu 2+ … Copper (II) Ion
