1. Chapter 4

2. 4.2 Symbols for the Elements 4.3 Dalton’s Atomic Theory
4.4 Formulas of Compounds
4.5 The Structure of the Atom
4.6 Introduction to the Modern Concept of Atomic Structure
4.7 Isotopes
4.8 Introduction to the Periodic Table
4.9 Natural States of the Elements
4.10 Ions
4.11 Compounds That Contain Ions
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3. Substances that cannot be broken down by simple chemical means
115 known: 88 found in nature, others are man made.
Just as you had to learn the 26 letters of the alphabet before you learned to read and write, you need to learn the names and symbols of the chemical elements before you can read and write chemistry.
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5. How the Term Element is Used
Could mean a single atom of that element (Ar or H).
Could mean molecules of an element (H2), which is hydrogen found in its natural state.
Could mean atoms of elements are present in some form (sodium found in the human body).
Look at each particular case to determine its proper use.
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6. Each element has a unique one- or two-letter symbol.
First letter is always capitalized and the second is not.
The symbol usually consists of the first one or two letters of the element’s name.
Examples: Oxygen O Krypton Kr
Sometimes the symbol is taken from the element’s original Latin or Greek name.
Examples: Gold Au aurum Lead Pb plumbum
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7. Names and Symbols of the Most Common Elements
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8. Most natural materials are mixtures of pure substances.
Pure substances are either elements or combinations of elements called compounds.
A given compound always contains the same proportions (by mass) of the elements.
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9. Law of Constant Composition
A given compound always has the same composition, regardless of where it comes from.
Water always contains 8 g of oxygen for every 1 g of hydrogen.
Carbon dioxide always contains 2.7 g of oxygen for every 1 g of carbon.
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10. Dalton’s Atomic Theory (1808)
Elements are made of tiny particles called atoms.
All atoms of a given element are identical.
The atoms of a given element are different from those of any other element.
Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms.
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11. Dalton’s Atomic Theory (continued)
Atoms are indivisible in chemical processes. Atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.
ONLY NUCLEAR REACTIONS CAN TRANSFORM ONE KIND OF ELEMENT TO ANOTHER
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12. Concept Check
Which of the following statements regarding Dalton’s atomic theory are still believed to be true?
Elements are made of tiny particles called atoms.
All atoms of a given element are identical.
A given compound always has the same relative
numbers and types of atoms.
IV. Atoms are indestructible.
Statements I and III are true. Statement II is not true (due to isotopes and ions). Statement IV is not true (due to nuclear chemistry).
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13. Chemical Formulas Describe Compounds
Compound – distinct substance that is composed of the atoms of two or more elements and always contains exactly the same relative masses of those elements.
Chemical Formulas – expresses the types of atoms and the number of each type in each unit (molecule) of a given compound.
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14. Rules for Writing Formulas
Each atom present is represented by its element symbol.
The number of each type of atom is indicated by a subscript written to the right of the element symbol.
When only one atom of a given type is present, the subscript 1 is not written.
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15. Exercise
The pesticide known as DDT paralyzes insects by binding to their nerve cells, leading to uncontrolled firing of the nerves. Before most uses of DDT were banned in the U.S., many insects had developed a resistance to it. Write out the formula for DDT. It contains 14 carbon atoms, 9 hydrogen atoms, and 5 atoms of chlorine.
C14H9Cl5
C14H9Cl5
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16. Postulated the existence of electrons using cathode-ray tubes.
J. J. Thomson (1898—1903)
Postulated the existence of electrons using cathode-ray tubes.
The atom must also contain positive particles that balance exactly the negative charge carried by particles that we now call electrons.
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17. J.J. Thomson, measured mass/charge of e-
Cathode-Ray Tube
J.J. Thomson, measured mass/charge of e-
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18. William Thomson (Plum Pudding Model)
Reasoned that the atom might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to counterbalance that positive charge.
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19. Explained the nuclear atom.
Ernest Rutherford (1911)
Explained the nuclear atom.
Atom has a dense center of positive charge called the nucleus.
Electrons travel around the nucleus at a relatively large distance.
A proton has the same magnitude of charge as the electron, but its charge is positive.
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20. Rutherford’s alpha particle scattering experiment.
5.5
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21. Rutherford’s alpha particle scattering experiment.
Deflection and scattering of alpha particles by positive gold nuclei.
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22. Rutherford and Chadwick (1932)
Most nuclei also contain a neutral particle called the neutron.
A neutron is slightly more massive than a proton but has no charge.
α particle is a helium nucleus:
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23. The nucleus contains protons and neutrons
We know atoms are composed of three main pieces - protons, neutrons and electrons
The nucleus contains protons and neutrons
The nucleus is only about cm in diameter
The electrons move outside the nucleus with an average distance of about 10-8 cm
therefore the radius of the atom is about 100,000 times larger than the radius of the nucleus
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24. Electrons – found outside the nucleus; negatively charged
The atom contains:
Electrons – found outside the nucleus; negatively charged
Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge
Neutrons – found in the nucleus; no charge; virtually same mass as a proton
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25. nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
The nucleus is:
Small compared with the overall size of the atom.
Extremely dense; accounts for almost all of the atom’s mass.
atomic radius ~ 100 pm = 1 x m
nuclear radius ~ 5 x 10-3 pm = 5 x m
“If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”
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27. Why do different atoms have different chemical properties?
The chemistry of an atom arises from its electrons.
Electrons are the parts of atoms that “intermingle” when atoms combine to form molecules.
It is the number of electrons that really determines chemical behavior.
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28. Isotopes
Atoms with the same number of protons but different numbers of neutrons.
Show almost identical chemical properties; chemistry of atom is due to its electrons.
In nature most elements contain mixtures of isotopes.
Mass Number X A Z
Element Symbol
Atomic Number H 1
H (D) 2
H (T) 3 U
235 92
238
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29. A – Z = n (number of neutrons)
Two Isotopes of Sodium
A – Z = n (number of neutrons)
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30. A – Z = n (number of neutrons)
Isotopes
X = the symbol of the element
Z = the atomic number (# of protons)
A = the mass number (# of protons and neutrons)
A – Z = n (number of neutrons)
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31. A – Z = n (number of neutrons)
Isotopes – An Example
C = the symbol for carbon
6 = the atomic number (6 protons)
14 = the mass number (6 protons and 8 neutrons)
C = the symbol for carbon
6 = the atomic number (6 protons)
12 = the mass number (6 protons and 6 neutrons)
A – Z = n (number of neutrons)
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32. What is the mass number of this isotope? Identify the element.
Exercise
A certain isotope X contains 23 protons and 28 neutrons.
What is the mass number of this isotope?
Identify the element.
Mass Number = 51
Vanadium
The mass number is 51. Mass Number = # protons + # neutrons. Mass Number = = 51. The element is vanadium. The number of protons determines the identity of the element.
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33. The Periodic Table
The periodic table shows all of the known elements in order of increasing atomic number.
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35. order elements by atomic mass saw a repeating pattern of properties
Periodic Law – When the elements are arranged in order of increasing relative mass, certain sets of properties recur periodically
used pattern to predict properties of undiscovered elements
where atomic mass order did not fit other properties, he re-ordered by other properties
Te & I
Mendeleev
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36. Periodic Pattern H Li Be B C N O F Na Mg Al Si P S Cl nm H2O a/b 1 H2
m Li2O b
7 LiH Be
m/nm BeO
a/b
9 BeH2
nm B2O3 a
11 ( BH3)n B
nm CO2 a
12 CH4 C
nm N2O5 a
14 NH3 N
nm O2
H2O O
nm OF2 HF F Na
m Na2O b
23 NaH
m MgO b
24 MgH2 Mg
m Al2O3
a/b
27 (AlH3) Al
nm/m SiO2 a
28 SiH4 Si
nm P4O10 a
PH3 P
nm SO3 a
32 H2S S
nm Cl2O7 a
35.5 HCl Cl
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37. Mendeleev's Predictions for Ekasilicon (Germanium)
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38. Periods – horizontal rows of elements
The Periodic Table
Metals vs. Nonmetals
Groups or Families – elements in the same vertical columns; have similar chemical properties
Periods – horizontal rows of elements
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39. Most elements are metals and occur on the left side.
The Periodic Table
Most elements are metals and occur on the left side.
The nonmetals appear on the right side.
Metalloids are elements that have some metallic and some nonmetallic properties.
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40. = Alkali Metals = Alkali Earth Metals = Noble Gases = Halogens
= Lanthanides
= Actinides
= Transition Metals
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41. Physical Properties of Metals
Section 4.9
Natural States of the Elements
Physical Properties of Metals
Efficient conduction of heat and electricity
Malleability (they can be hammered into thin sheets)
Ductility (they can be pulled into wires)
A lustrous (shiny) appearance
High densities
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42. Most elements are very reactive.
Natural States of the Elements
Section 4.9
Most elements are very reactive.
Elements are not generally found in uncombined form.
Exceptions are:
Noble metals – gold, platinum and silver
Noble gases – Group 8
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43. Natural States of the Elements
Section 4.9
Diatomic Molecules
Nitrogen gas contains N2 molecules.
Oxygen gas contains O2 molecules.
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44. Natural States of the Elements
Section 4.9
Molecular Elements
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45. Natural States of the Elements
Section 4.9
Allotropes
Different forms of a given element.
Example:
Solid carbon occurs in three forms.
Diamond
Graphite
Buckminsterfullerene
Carbon Allotropes
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46. Atoms can form ions by gaining or losing electrons.
Section 4.10
Ions
Atoms can form ions by gaining or losing electrons.
Metals tend to lose one or more electrons to form positive ions called cations.
Cations are generally named by using the name of the parent atom.
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47. Ions
Section 4.10
Nonmetals tend to gain one or more electrons to form negative ions called anions.
Anions are named by using the root of the atom name followed by the suffix –ide.
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48. Ion Charges and the Periodic Table
Ions
Section 4.10
Ion Charges and the Periodic Table
The ion that a particular atom will form can be predicted from the periodic table.
Group or Family
Charge
Alkali Metals (1A) 1+
Alkaline Earth Metals (2A) 2+
Halogens (7A) 1–
Noble Gases (8A)
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49. Ion Charges and the Periodic Table
Section 4.10
Ions
Ion Charges and the Periodic Table
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50. An ion with a 3+ charge contains 23 electrons. Which ion is it?
Section 4.10
Ions
Exercise
An ion with a 3+ charge contains 23 electrons. Which ion is it?
a) Fe3+
b) V3+
c) Ca3+
d) Sc3+
+ve charge ≡ electrons are lost
-ve charge ≡ electrons are gained
# of e- = Z – positive charge or
# of e- = Z + negative charge
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51. A – Z = n (number of neutrons)
Section 4.10
Ions
A – Z = n (number of neutrons)
Exercise
A certain ion X1+ contains 54 electrons and 78 neutrons.
What is the mass number of this ion?
133
Z = # of e- + positive charge or
Z = # of e- - negative charge
+ve charge ≡ electrons are lost
-ve charge ≡ electrons are gained
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52. Ions combine to form ionic compounds. Properties of ionic compounds
Section 4.11
Compounds That Contain Ions
Ions combine to form ionic compounds.
Properties of ionic compounds
High melting points
Conduct electricity
If melted
If dissolved in water
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53. Ionic compounds are electrically neutral.
Section 4.11
Compounds That Contain Ions
Ionic compounds are electrically neutral.
The charges on the anions and cations in the compound must sum to zero.
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54. Formulas for Ionic Compounds
Section 4.11
Compounds That Contain Ions
Formulas for Ionic Compounds
Write the cation element symbol followed by the anion element symbol.
The number of cations and anions must be correct for their charges to sum to zero.
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55. Z = # of e- + positive charge or Z = # of e- - negative charge
Section 4.11
Compounds That Contain Ions
Concept Check
A compound contains an unknown ion X and has the formula XCl2. Ion X contains 20 electrons. What is the identity of X?
a) Ti2+
b) Sc+
c) Ca2+
d) Cr2+
Z = # of e- + positive charge or
Z = # of e- - negative charge
The charge on the cation must be +2 since there are two Cl’s each with a –1 charge (giving an overall charge of –2 for the anion side). The cation can now be represented as X2+, containing 20 electrons. Therefore, 22 protons must be present to give a charge of +2 (+22p – 20e = +2). The element with 22 protons is titanium.
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56. a) CaBr2 b) KrBr c) RbBr d) SrBr2 Compounds That Contain Ions
Section 4.11
Compounds That Contain Ions
Concept Check
A member of the alkaline earth metal family whose most stable ion contains 36 electrons forms a compound with bromine. What is the correct formula for this compound?
a) CaBr2
b) KrBr
c) RbBr
d) SrBr2
Ionic compounds are electrically neutral.
The charges on the anions and cations in the compound must sum to zero.
An element in the alkaline earth metal family (Group 2) forms a +2 charge when forming a compound. Therefore two Br– ions will be required to give a net zero charge overall. 38 protons must be present to give a charge of +2 (+38p – 36e = +2). The element with 38 protons is strontium. The compound is therefore SrBr2.
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57. Homework Problems: pages 107 - 111
Chapter 4 Homework
Homework
Reading assignment
Pages 74 through 105
Homework Problems: pages
Questions and problems 3, 5, 9, 13, 19, 25, 27, 31, 33, 35, 37, 39, 42, 45, 47, 49, 51, 53, 57, 59, 61, 63, 67, 69, 71, 73, 75, 77, 81.
Due on
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