1. Chapter 2 Atoms, Molecules and Ions

2. Atomic Theory of Matter

3. Law of Conservation of Mass
The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

4. Law of Definite ( or constant ) composition
No matter what its source, a particular chemical compound is composed of the same elements in the same proportions by mass.

5. Radioactivity
Radioactivity is the spontaneous emission of radiation by an atom. It was first observed by Henri Becquerel. Marie and Pierre Curie also studied it.
Three types of radiation were discovered by Ernest Rutherford:
 particles (a helium ion with a +2 charge)
 particles(high-energy, high-speed electron or positron)
 rays

6. The atom Protons were discovered by Rutherford in 1919.
Atomic mass unit (symbolized AMU or amu) is defined as precisely 1/12 the mass of an atom of carbon-12. = = x g
Protons were discovered by Rutherford in 1919.
(charge = +1 = x C)
 Mass :  a.m.u.   or  x 10-24 g.
Proton is 1837 times heavier than an electron.
Neutrons were discovered by James Chadwick in (neutral)
Mass = a.m.u. or  x 10-24 g.
Neutron is 1842 times heavier than an electron.
Electrons: discovered by J. J. Thomson (1897) as Streams of negatively charged particles that emanate from cathode tubes
Charge of electron = -1 = x C
Mass of electron = x 10-4 (amu) = x g

7. Symbols of Elements Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons

8. Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei (have different masses)
Mass Number X A Z
Element Symbol
Atomic Number H 1
H (D) 2
H (T) 3 U
235 92
238

9. Atomic and Molecular Masses
The atomic mass (ma) is the mass of an atom in (amu). For atoms, the protons and neutrons of the nucleus account for almost all of the mass, and the atomic mass has nearly the same value as the mass number.
The  molecular mass is the mass of a molecule. It is calculated as the sum of the atomic weights of each constituent element multiplied by the number of atoms of that element in the molecular formula.
* Atomic and molecular masses can be measured with great accuracy using a mass spectrometer.
Formula weights are generally reported for ionic compounds.

10. Average Mass

11. Periodic Table
The periodic table is a systematic catalog of the elements.
Elements are arranged in order of atomic number.

12. Periodic Table The rows on the periodic chart are periods.
Columns are groups.
Elements in the same group have similar chemical properties.

13. Periodicity
Periods: When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.
Groups:Five groups are known by their names.

14. Nonmetals are on the right side of the periodic table (with the exception of H).
Metalloids border the stair-step line (with the exception of Al, Po, and At
Metals are on the left side of the chart.

15. Chemical Formulas
The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.
Molecular compounds are composed of molecules and almost always contain only nonmetals.

16. Diatomic Molecules
These seven elements occur naturally as molecules containing two atoms:
Hydrogen
Nitrogen
Oxygen
Fluorine
Chlorine
Bromine
Iodine

17. Types of Formulas
Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.
Molecular formulas give the exact number of atoms of each element in a compound.
Structural formulas show the order in which atoms are bonded.
Perspective drawings also show the three-dimensional array of atoms in a compound.

18. H2O
molecular
empirical
H2O
C6H12O6
CH2O O3 O
N2H4
NH2

20. Ions When atoms lose or gain electrons, they become ions.
Cations are positive and are formed by elements on the left side of the periodic chart.
Anions are negative and are formed by elements on the right side of the periodic chart.

21. cation – ion with a positive charge
An ion is an atom, or group of atoms, that has a net positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation. Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl-
17 protons
18 electrons Cl
17 protons
17 electrons

22. How many protons and electrons are in
Do You Understand Ions?
How many protons and electrons are in Al 27 13 ? 3+
13 protons, 10 (13 – 3) electrons
How many protons and electrons are in Se 78 34 2- ?
34 protons, 36 (34 + 2) electrons

23. Solution Practice Exercise
Sample Exercise 2.7 Writing Chemical Symbols for Ions
Give the chemical symbol, including superscript indicating mass number, for (a) the ion with 22 protons, 26 neutrons, and 19 electrons; (b) the ion of sulfur that has 16 neutrons and 18 electrons.
Solution
(a) The number of protons is the atomic number of the element. A periodic table or list of elements tells us that the element with atomic number 22 is titanium (Ti). The mass number (protons plus neutrons) of this isotope of titanium is = 48. Because the ion has three more protons than electrons, it has a net charge of 3+: 48Ti3+.
(b) The periodic table tells us that sulfur (S) has an atomic number of 16. Thus, each atom or ion of sulfur contains 16 protons. We are told that the ion also has 16 neutrons, meaning the mass number is = 32. Because the ion has 16 protons and 18 electrons, its net charge is 2– and the ion symbol is 32S2–.
In general, we will focus on the net charges of ions and ignore their mass numbers unless
the circumstances dictate that we specify a certain isotope.
Practice Exercise
How many protons, neutrons, and electrons does the 79Se2– ion possess?
Answer: 34 protons, 45 neutrons, and 36 electrons

24. Ionic Bonds
Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

25. Sample Exercise 2.9 Identifying Ionic and Molecular Compounds
Which of these compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4?
Solution
We predict that Na2O and CaCl2 are ionic compounds because they are composed of a metal combined with a nonmetal. We predict (correctly) that N2O and SF4 are molecular compounds because they are composed entirely of nonmetals.
Practice Exercise
Which of these compounds are molecular: CBr4, FeS, P4O6, PbF2?
Answer: CBr4 and P4O6 25

26. Writing Formulas
Because compounds are electrically neutral, one can determine the formula of a compound this way:
The charge on the cation becomes the subscript on the anion.
The charge on the anion becomes the subscript on the cation.
If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.

27. The ionic compound NaCl
ionic compounds consist of a combination of cations and an anions
they only have empirical formulas, no molecular formula
the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero
The ionic compound NaCl

28. Formula of Ionic Compounds
2 x +3 = +6
3 x -2 = -6
Al2O3
Al3+
O2-
1 x +2 = +2
2 x -1 = -2
CaBr2
Ca2+
Br-
1 x +2 = +2
1 x -2 = -2
Na2CO3
Na+
CO32-

29. Common Cations

30. Common Anions

31. Inorganic Nomenclature
Write the name of the cation.
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.
If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

32. Patterns in Oxyanion Nomenclature
When there are two oxyanions involving the same element:
The one with fewer oxygens ends in -ite.
The one with more oxygens ends in -ate.
NO2− : nitrite; SO32− : sulfite
NO3− : nitrate; SO42− : sulfate

33. Central atoms on the second row have bond to at most three oxygens; those on the third row take up to four.
Charges increase as you go from right to left.

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Acids
HClO HClO HClO HClO
perchloric acid chloric acid chlorous acid hypochlorous acid
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35. Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi- (one H), dihydrogen (two H) etc., to the name as follows:
• CO32– is the carbonate anion.
• HCO3– is the hydrogen carbonate
(or bicarbonate) anion.
• PO43– is the phosphate ion.
• H2PO4– is the dihydrogen phosphate anion.
HPO42– is the hydrogen phosphate anion.

36. Ionic Compounds (salts)
often a metal + nonmetal (or polyatomic ion)
anion (nonmetal), add “ide” to element name
BaCl2
barium chloride
K2O
potassium oxide
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate

37. Transition metal ionic compounds
indicate charge on metal with Roman numerals
FeCl2
iron(II) chloride
2 Cl- -2 so Fe is +2
FeCl3
3 Cl- -3 so Fe is +3
iron(III) chloride
Cr2S3
3 S-2 -6 so Cr is +3 (6/2)
chromium(III) sulfide

38. An acid can be defined as a substance that yields
hydrogen ions (H+) when dissolved in water.
HCl Pure substance, hydrogen chloride
Dissolved in water (H+ Cl-), hydrochloric acid
HCl: hydrochloric acid
HBr: hydrobromic acid
HI: hydroiodic acid
An oxoacid is an acid that contains hydrogen, oxygen, and another element.
HNO3
nitric acid
H2CO3
carbonic acid
H2SO4
sulfuric acid

39. Solution Practice Exercise
Sample Exercise 2.14 Relating the Names and Formulas of Acids
Name the acids (a) HCN, (b) HNO3, (c) H2SO4, (d) H2SO3.
Solution
(a)The anion from which this acid is derived is CN–, the cyanide ion. Because this ion has an -ide ending, the acid is given a hydro- prefix and an -ic ending: hydrocyanic acid. Only water solutions of HCN are referred to as hydrocyanic acid. The pure compound, which is a gas under normal conditions, is called hydrogen cyanide. Both hydrocyanic acid and hydrogen cyanide are extremely toxic.
(b) Because NO3– is the nitrate ion, HNO3 is called nitric acid (the -ate ending of the anion is replaced with an -ic ending in naming the acid).
(c)Because SO42– is the sulfate ion, H2SO4 is called sulfuric acid.
(d) Because SO32– is the sulfite ion, H2SO3 is sulfurous acid (the -ite ending of the anion is replaced with an -ous ending).
Practice Exercise
Give the chemical formulas for (a) hydrobromic acid, (b) carbonic acid.
Answer: (a) HBr, (b) H2CO3 39

40. Nomenclature of Binary Compounds
The less electronegative atom is usually listed first.
A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however) .

41. Nomenclature of Binary Compounds
The ending on the more electronegative element is changed to -ide.
CO2: carbon dioxide
CCl4: carbon tetrachloride

42. Nomenclature of Binary Compounds
If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one.
N2O5: dinitrogen pentoxide

43. Molecular Compounds HI hydrogen iodide NF3 nitrogen trifluoride SO2
sulfur dioxide
N2Cl4
dinitrogen tetrachloride
NO2
nitrogen dioxide
N2O
dinitrogen monoxide

44. Solution Practice Exercise
Sample Exercise 2.15 Relating the Names and Formulas of Binary Molecular Compounds
Name the compounds (a) SO2, (b) PCl5, (c) Cl2O3.
Solution
The compounds consist entirely of nonmetals, so they are molecular rather than ionic. Using the prefixes in Table 2.6,
we have (a) sulfur dioxide, (b) phosphorus pentachloride,
and (c) dichlorine trioxide.
Practice Exercise
Give the chemical formulas for (a) silicon tetrabromide,
(b) disulfur dichloride.
Answer: (a) SiBr4, (b) S2Cl2 44

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