Download Chapter 2 Atoms, Molecules, and Ions

Chapter 2 Atoms, Molecules, and Ions
2.1 The Atomic Theory of Matter
Is there such thing like a indivisible particle of matter (atom)?
There was no evidence, so people could not decide one way or another.
Based on the following observations:
1.
Law of conservation of mass/matter (Antoine Lavoisier). (Total mass before rxn. = total mass after rxn.)
2.
Law of constant composition/definite proportions. (Joseph Proust) A given compound always contains the
same elements in the same ratio.
Proportions in Sodium Chloride
A 100.0 g sample of sodium chloride contains 39.3 g of sodium and 60.7 g of chlorine
A 200.0 g sample of sodium chloride contains 78.6 g of sodium and 121.4 g of chlorine
A 58.44 g sample of sodium chloride contains 22.99 g of sodium and 35.44 g of chlorine
Dalton’s Atomic Theory (1803)
1. Elements are composed of atoms: indivisible, indestructible.
2. The atoms of a given element are similar to each other and different from the atoms of all other elements.
3. Chemical compounds consist of atoms combined in specific, whole number, ratios. (Constant Composition)
4. Chemical reactions change only the way the atoms are combined in compounds; the atoms themselves are unchanged.
(Conservation of Mass)
What Makes a Theory (Model) a Good Theory?
A good theory (model):
a. explains know facts.
b. predicts new ones…
a. Dalton’s theory explained the laws of conservation of matter and constant composition. (How?)
b. Dalton predicted the law of multiple proportions - two elements can form more than one compound by combining in
different ratios. (This was later confirmed.)
H2O and H2O2;
CO and CO2
2.2 The Discovery of Atomic Structure
Thomson’s Experiment
J. J. Thompson,1897. (Cathode rays = beam of electrons) Describe the experiment.
In the presence of electric and magnetic forces, when the two forces are balanced, the charge to mass ratio can be calculated
(see the handout to look at the math) and Thompson found it to be:
charge/mass ratio = 1.76  108 coulombs/g.
Robert Millikan determined the charge on the electron in 1909. (So much patience!)
1. He watched drops of oil fall due to gravity (the mass was determined).
2. The oil drops were irradiated with X-rays and as a result became charged.
3. When the electrical force on the charged drops balanced the gravity force, then oil drops were suspended on air.
Different droplets had different charges, but all were a multiple of: 1.602  10-19 coulombs.
The conclusion was:
• Charge of electron: 1.602  10-19 coulombs
• Mass of electron: 9.10  10-28 g (from Thompson’s ratio.)
1
Fossum-Reyes
Radioactivity (1896): the spontaneous emission of radiation by an atom.
Three types of radiation were discovered by Ernest Rutherford:
–  particles
–  particles
–  rays
Atomic Structure
By the early 1900’s scientists already new atoms contain electrons (negatively charged), but they are neutral overall… (i.e.
they must contain something that neutralizes the charge).
Thomson’s Model
J.J. Thomson proposed the Plum Pudding Model (didn’t last long…): He thought of the atom as a sphere of uniformly
distributed positive charge with electrons (negatively charged) embedded in it (like a chocolate chip cookie) to yield a neutral
atom.
Ernest Rutherford’s Experiment (1911)
Shot α particles (high energy, positively charged and very massive) through a thin metal foil. If the plum pudding model was
correct, the α particles should pass straight through the foil with a few minor deflections at most (why?). What he found
instead, was that some α particles suffered large deflections and some even bounced backwards, although most did go
straight through. (He described the phenomenon as “shooting a cannonball through tissue paper and having the ball
bouncing back”.)
Conclusion: atoms contain a very tiny central core (the nucleus) that contains all of the positive charge and most of the mass
of the atom. The atom is mostly empty space.
2.3 Modern view of Atomic Structure
• Protons were discovered by Rutherford in 1919.
• Neutrons were discovered by James Chadwick in 1932.
The Nuclear Atom
Particle
Charge
Charge (Atomic
Mass
Mass (Atomic
Relative Mass
Scale)
Scale)
Proton (p+)
+ 1.602 x 10-19 C
+1
1.673 x 10-27 kg
1.0073
≈1
0
Neutron (n )
0
0
1.675 x 10-27 kg
1.0087
≈1
-19
-31
-4
Electron (e )
– 1.602 x 10 C
–1
9.1 x 10 kg
5.486 x 10
≈0
If the atom is compared to a baseball stadium, the nucleus would be the size of a pea; nevertheless, the mass of that peasized nucleus would be… 250 million tons! The nucleus is extremely dense.
Atoms have no net electrical charge.
Atomic Number, Mass Number and Isotopes
Atomic number (Z) = # p+
Mass number (A) = # p+ + # n0
*In a neutral atom, the number of e- and p+ must be equal.*
Isotopes
•
Isotopes are atoms of the same element with different masses (i.e. they have different numbers of neutrons).
11
12
13
;
; 146𝐶
6𝐶 ;
6𝐶
6𝐶
Isotope symbol
𝑀𝑎𝑠𝑠 𝑁𝑢𝑚𝑏𝑒𝑟 𝐴
𝐴𝑡𝑜𝑚𝑖𝑐 𝑁𝑢𝑚𝑏𝑒𝑟 𝑍𝑋
(𝐸𝑙𝑒𝑚𝑒𝑛𝑡 ′ 𝑠 𝑠𝑦𝑚𝑏𝑜𝑙)
For Carbon -12 → 126𝐶
𝑜𝑟
12
𝐶 (𝑠𝑜𝑚𝑒𝑡𝑖𝑚𝑒𝑠 𝑍 𝑖𝑠 𝑜𝑚𝑖𝑡𝑡𝑒𝑑)
2
Fossum-Reyes
Problem 1
Q. Give the number of each type of particle (p+, n0 and e-) for the following isotopes.
a.
32
15
b.
18
8
P
O
Q. Write the isotope symbol for an atom of nickel (Z = 28) containing 32 neutrons.
2.4 Atomic Weights
Atomic Mass Units (amu) are a more convenient unit:
1 amu = 1.66054 x 10-24 g
An atom of C-12 weighs exactly 12 amu. All other masses are measured relative to C-12.
The Periodic Table Shows a Weighted Average
𝐴𝑡𝑜𝑚𝑖𝑐 𝑊𝑒𝑖𝑔ℎ𝑡 = ∑(𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑎𝑠𝑠) × (𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛𝑎𝑙 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒)
𝐴𝑡𝑜𝑚𝑖𝑐 𝑊𝑒𝑖𝑔ℎ𝑡 = (35 𝑎𝑚𝑢 × 0.7578) + (37 𝑎𝑚𝑢 × 0.2422)
𝐴𝑡𝑜𝑚𝑖𝑐 𝑊𝑒𝑖𝑔ℎ𝑡 = 35.45 𝑎𝑚𝑢
Problem 2
Calculate the weighed-average atomic mass of Magnesium.
Mg
24
23.9850 amu 78.99%
25
24.9858 amu 10.00%
26
25.9826 amu 11.01%
100.00%
Mass spectrometer (p.49)
Separates samples in fractions of different mass; the fractional abundances are indicated by the relative signal intensities.
2.5 The Periodic Table
When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. (Periodicity)
Metals are on the left side of the chart. Solids at room temperature, except Hg..
1. Reflective surface (shiny).
2. Conduct heat and electricity.
3. Malleable (can be shaped) and ductile (pulled into wires).
4. Form cations (lose electrons to form positive ions).
Nonmetals are on the right side of the periodic table (with the exception of H). Found in all three states.
1. Poor conductors of heat and electricity.
2. Their solids are brittle.
3. Form anions (gain electrons to form negative ions).
4. Found in all three states.
3
Fossum-Reyes
Metalloids border the stair-step line (with the exception of Al, Po, and At) – intermediate properties.
1. Show some properties of metals and some of nonmetals.
2. Also known as semiconductors.
2.6 Molecules and Molecular Compounds
Molecules – two or more atoms joined together.
Molecular compounds (composed of molecules) almost always contain only nonmetals.
Chemical Formulas indicate the type and number of atoms found (subscripts).
•
•
Certain elements occur as 2 atom molecules
 rule of 7’s (Memorize them!)
Other elements occur as polyatomic molecules
 P4, S8, Se8
Types of Formulas
• Empirical formulas - lowest whole-number ratio of atoms.
• Molecular formulas - exact number of atoms.
Empirical Formula
Molecular Formula
CH
C6H6 (Benzene)
CH
C2H2 (Ethene)
•
•
Structural formulas – how atoms are bonded.
3D formulas (Perspective drawings) - show the three-dimensional array of atoms.
Molecular and Structural Formulas
Molecular Formula
Structural Formula
O
H2O
H
O
CO2
C2H6O
H
H
C
H
H
C
C
O
H
H
H
C
H
CH3CH2OH
H
H
C2H6O
O
H
O
C
H
CH3OCH3
H
2.7 Ions and Ionic Compounds
By gaining or losing one or more electrons neutral atoms become ions.
• Cations (metals) are positive
• Anions (nonmetals) are negative
Ion Charge and the Periodic Table
• The charge on an ion can often be determined from an element’s position on the Periodic Table
• Metals always form positively charged cations
• For many main group metals, the charge = the group number
• Nonmetals form negatively charged anions
• For nonmetals, the charge = the group number − 8
4
Fossum-Reyes
Elements try to lose or gain electrons to end up with the same number of electrons as the noble gas closest to them (octet
rule).
Practice – What is the charge on each of the following ions?
• potassium cation
• sulfide anion
• calcium cation
• bromide anion
• aluminum cation
Ionic Compounds
Ionic compounds (such as NaCl) are generally formed between metals and nonmetals (ions).
Formulas of ionic compounds
The rule is: Compounds are electrically neutral.
Ions Present
Formula
If the subscripts are not the lowest whole-number ratio, divide them by the greatest common factor.
Formulas of Ionic Compounds
Ions
Formula
Li+ and N3Ca2+ and ClNa+ and S2Al3+ and BrProblem 3
Use the ions of the pairs of elements below to give the formula for their ionic compound.
a. Ba and O
b. Al and S
c. K and P
Names of Common Ions
Memorize: Their Structure, Charge and Name. (See the handout on common ions.) Look for patterns.
Polyatomic Ions
2.8 Naming Compounds
1. Classify the compound:
a) Ionic – metal and nonmetal (contains ions).
b) Acid – formula begins with H.
c) Binary molecular – two nonmetals.
d) Organic – contains C, H, possibly O, N (carbon based).
2. Use appropriate rules to name the compound.
5
Fossum-Reyes
Naming Ionic Compounds
a. Write the name of the cation. (Name of the element).
b. If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the
polyatomic ion.
c. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.
d. Name the cation, then the anion. (We do not need to specify how many of each there are. Why?)
Naming Ionic Compounds
Compound
Contains
BaBr2
Ba2+, Br–
KOH
K+,OH–
SnCl2
Sn2+, Cl–
SnCl4
Sn4+, Cl–
PbO2
Pb4+, O2–
MgSO4
Mg2+, SO42–
Fe3(PO4)2
Fe2+, PO43–
FePO4
Fe3+, PO43–
(NH4)2CO3
NH4+, CO32–
Name
Barium bromide
Potassium hydroxide
Tin (II) chloride
Tin (IV) chloride
Lead (IV) oxide
Magnesium sulfate
Iron (II) phosphate
Iron (III) phosphate
Ammonium carbonate
Problem 4
Name:
CaI2
Problem 5
Write the formulas:
Calcium hypochlorite
K2SO4
Sodium bicarbonate
(NH4)2S
Cobalt (III) oxide
CuF
Lithium phosphate
CuF2
Nickel (II) cyanide
Ag2CO3
Aluminum nitrite
Cr(NO3)3
Naming Acids – Contain H+ and an anion.
a. Binary acids (its anion ends in “_ide”)
Hydro _____ic acid
HCl
Hydrochloric acid
HBr
Hydrobromic acid
H2S
Hydrosulfuric acid
b. Oxyacids (anion ends in “_ite” or “_ate”.)
The name is related to the name of the anion:
If the anion ends in:
acid ends in:
“_ate”
“_ic acid”
“_ite”
“_ous acid”
HNO3
HNO2
H2SO4
H2SO3
H3PO4
NO3–
NO2–
SO4 2 –
SO3 2 –
PO4 3 –
nitrate
nitrite
sulfate
sulfite
phosphate
nitric acid
nitrous acid
sulfuric acid
sulfurous acid
phosphoric acid
6
Fossum-Reyes
Problem 6
Name:
HC2H3O2
;
HI
H2CO3
;
HClO4
Problem 7
Give the Formula:
Hydrofluoric acid
;
Chlorous acid
Hypochlorous acid
;
Chromic acid
Naming Binary Molecular Compounds
Made of two nonmetals, these compounds DO NOT contain ions.
Common names:
H2O Water NH3 Ammonia PH3 Phosphine
Naming Binary Molecular Compounds
Systematic names use numeral prefixes
(since they do not have charges, we
can’t predict how many atoms of each
element are present):
- Name the first element.
- Second element ends in “_ide”.
- Don’t use “mono” for the first element.
Naming Binary Molecular Compounds
CO2
carbon dioxide
CO
carbon monoxide
(just one “o”)
N2O
dinitrogen monoxide
PCl5
phosphorus pentachloride
N2O4
dinitrogen tetroxide
(“a” omitted)
Naming Binary Molecular Compounds
Note:
Pure Substance
Aqueous Solution
HCl(g)
hydrogen chloride
hydrochloric acid
H2S(g)
hydrogen sulfide
hydrosulfuric acid
Problem 8
Name:
ClF3
;
Formulas:
Phosphorus triiodide
P4O10
;
Sulfur dichloride
SO3
;
Xenon trioxide
Problem 9
Name:
Co(NO2)2
;
PCl3
HBrO3
;
CaSO3
N2F4
;
Zn(C2H3O4) 2
HF
7
Fossum-Reyes
Problem 10
Formulas:
Tin (IV) permanganate
;
Sulfurous acid
Lithium sulfite
;
Dioxygen difluoride
Silver oxide
;
Periodic acid
Nickel (II) carbonate
2.9 Organic Compounds
Organic chemistry is the study of carbon; it has its own system of nomenclature. There are literally millions of organic
compounds, which are divided in families (posses certain structural features that identify them).
Alkanes contain C and H only and no multiple bonds (the simplest hydrocarbons).
H
H
C
H
H
H
CH4
Methane
H
H
C
C
H
H
H
H
C2H6
Ethane
H
H
H
C
C
C
H
H
H
H
C3H8
Propane
The family of alcohols: molecules that contain the –OH group. We use the –ol ending.
H
H
C
O
H
CH3OH
Methanol
H
H
H
H
C
C
H
H
C2H5OH
Ethanol
O
H
H
H
H
H
C
C
C
H
H
H
O
H
C3H7OH
1-Propanol
8
Fossum-Reyes