1. what is an atom made of (think of what you have learned in other classes)? day 3 10-16
TRANSCRIPT
1. What is an atom made of (think of what you have learned in other classes)?
Day 3 10-16
Atoms, Molecules, and Ions
Chapter 2
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
1. Brainstorm with your discussion partner: everything you know about an atom.
2. Look up the definition in the back of the book, and record it
3. Draw an atom of hydrogen. Draw an atom of oxygen. (There is no wrong answer at this time… Draw something!)
ATOM Neutron
Proton
Electron
(0)
(+)
(-)
Nucleus
When and Why do we use models?
In chemistry:
I. Early Greeks• Everything is made
of 4 elements: _______________, _______________,
_______________, and _______________.
• These combine and interact to make everything.
EarthEarthFireFireAirAir
WaterWater
Continuous matter – accepted for nearly 200 years
-Aristotle & Plato
400 B.C. Basic particle = an atom – “indivisible”
-Democritus
The Atom in Ancient Greece
Sparks of Knowledge - 1700’s
• Law of conservation of matter states that during any ________ or _________ process, matter is neither _______ nor ________
• In a reaction, the mass of the reactants ___________ the mass of the products.
chemicalchemical physicalphysicalcreatedcreated destroyeddestroyed
E Q U A L SE Q U A L S
Sparks of Knowledge - 1700’s Joseph Proust
• Developed the law of definite proportions or constant composition which
states that the _____ ratio of elements in a compound is always _________.
• Examples:
massmassthe samethe same
Water 1g H : 8g Water 1g H : 8g OOCarbon Dioxide 3g C : Carbon Dioxide 3g C : 8g O8g O
Law of Conservation of Mass – Mass is neither destroyed nor created
Reactantsmass = Productsmass
Law of Definite Proportions – A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample
Lasting Laws
Dalton’s Atomic Theory (1800s):
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements combine in simple whole-numbered ratios to form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.
Modern Atomic Theory:
1. All matter is composed of extremely small particles called atoms.
2. A GIVEN ELEMENT CAN HAVE ATOMS WITH DIFFERENT MASSES
3. ATOMS ARE DIVISIBLE INTO EVEN SMALLER PARTICLES.
4. Atoms of different elements combine in simple whole-numbered ratios to form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.
Dalton’s Atomic Theory (1800s):
1.All matter is composed of extremely small particles called atoms.
Modern Atomic Theory:
1.All matter is composed of extremely small particles called atoms.
Dalton’s Atomic Theory (1800s):
2.Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
Modern Atomic Theory:
2.A GIVEN ELEMENT CAN HAVE ATOMS WITH DIFFERENT MASSES
Dalton’s Atomic Theory (1800s):
3.Atoms cannot be subdivided, created, or destroyed.
Modern Atomic Theory:
3.ATOMS ARE DIVISIBLE INTO EVEN SMALLER PARTICLES.
Dalton’s Atomic Theory (1800s):
4.Atoms of different elements combine in simple whole-numbered ratios to form chemical compounds.
Modern Atomic Theory:
4.Atoms of different elements combine in simple whole-numbered ratios to form chemical coms.
Dalton’s Atomic Theory (1800s):
5. In chemical reactions, atoms are combined, separated, or rearranged.
Modern Atomic Theory:
5. In chemical reactions, atoms are combined, separated, or rearranged.
1. One thing Dalton had right was …
2. One thing Dalton was wrong about was…
Pd 3 Day 4 10-19
We have Atoms! – 1803-09 John Dalton
Atom’s Appearance• __________• __________• _____ ________• Different __________ • Different __________• Different __________
SolidSolidSphereSphereLike marblesLike marbles
sizessizescolorscolors
weightsweights
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Dalton’s Atomic Theory (1808)1. Elements are composed of extremely small particles
called atoms.
2. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements.
3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction.
4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.
Review sections 2.1 and 2.2 and complete #s 2.1-2.8 on page 68.
- due Tuesday
20
1. One thing Dalton had right was …
2. One thing Dalton was wrong about was…
Pd 1 Day 4 10-19
1. One thing Dalton had right was …
2. One thing Dalton was wrong about was…
Pd 8 Day 5 10-20
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8 X2Y16 X 8 Y+
Law of Conservation of Mass
We have Atoms! - 1850’s • Cathode Ray Tubes
–Invented and studied in the mid-1800’s
–A beam of cathode ray particles . . .
• applies ________,
• makes the temperature of objects _________,
• and is affected by ___________ and ___________ fields.
forceforce
warmerwarmerelectricelectric
magneticmagnetic
We have Atoms! - 1850’s Cathode Ray Tubes
We have Atoms! - 1850’s Cathode Ray Tubes
We have Atoms! – 1897 J. J. Thomson
• Studied _________ ______.
• All metals create _______________________.
• Cathode ray particles are _____ and _________ _________.
• Thomson coined the _______ as another name for ______ ___ _________.
cathode cathode raysrays
identical cathode identical cathode raysrays tinytiny
negatively negatively chargedcharged
electronelectron
particlesparticlescathode cathode ray ray
We have Atoms! – 1897 J. J. Thomson
solidsolidspheresphere
negative electronsnegative electronspositive positive surroundingssurroundings
• Atom’s appearance:
–______
–_______–Parts
•________________
•_______ ___________
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Thomson’s Model
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J.J. Thomson, measured mass/charge of e-
(1906 Nobel Prize in Physics)
Cathode Ray Tube
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Cathode Ray Tube
Review sections 2.1 and 2.2 and complete #s 2.1-2.8 on page 68.
- due Tuesday
32
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e- charge = -1.60 x 10-19 C
Thomson’s charge/mass of e- = -1.76 x 108 C/g
e- mass = 9.10 x 10-28 g
Measured mass of e-
(1923 Nobel Prize in Physics)
Millikan’s Experiment
Review sections 2.1 and 2.2 and complete #s 2.1-2.8 on page 68.
- due Tuesday
34
• Investigated _________ and its behaviors.
• The chemical properties of radioactive elements _______ as radiation is emitted.
radiationradiation
changechange
We have Atoms! – 1898Marie Skłodowska Curie and Pierre Curie
• Isolated and discovered the elements _________ and ________.
poloniumpoloniumradiumradium
We have Atoms! – 1898Marie Skłodowska Curie and Pierre Curie
We have Atoms! – 1899 Ernest Rutherford
• Analyzed ________ to determine its ___________.
–______________, , are fairly heavy with a ________ charge.
–_____________, , are very tiny with a ________ charge.
radiationradiationcompositioncompositionAlpha particlesAlpha particles
positivepositiveBeta particlesBeta particles
negativenegative
We have Atoms! – 1899 Ernest Rutherford
–_____________, , are very tiny with a ________ charge.
–____________, , have no mass with a _______ charge.
Beta particlesBeta particlesnegativenegative
Gamma raysGamma raysneutralneutral
Radiation Penetration
Radiation Separation
• What does this separation pattern indicate about charges and masses of these types of radiation?–α: ______ mass and ________ charge
–β: ________ mass and _________ charge
largerlarger positivepositivesmallersmaller negativenegative
• What does this separation pattern indicate about charges and masses of these types of radiation?–β: ________ mass and _________
charge
–γ: ___ mass and ___ charge
smallersmaller negativenegative
nono nono
Radiation Separation
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(uranium compound)
Types of Radioactivity
Atoms defy what we thought we knew! 1909-1910
• Ernest Rutherford’s–Experiment applet #1
–Experiment applet #2
• Ernest Rutherford’s
Atom’s appearance:– ____________________– ____________________________________– ________________________________
Hollow Hollow spheresphereDense center with positive chargeDense center with positive charge
Electrons occupy empty spaceElectrons occupy empty space
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1. atoms positive charge is concentrated in the nucleus2. proton (p) has opposite (+) charge of electron (-)3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)
particle velocity ~ 1.4 x 107 m/s(~5% speed of light)
(1908 Nobel Prize in Chemistry)
Rutherford’s Experiment
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atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
Rutherford’s Model of the Atom
“If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”
Atoms defy what we thought we knew! 1932
• James Chadwick discovered _________
–Have no _______
–Located in the atom’s
_______
neutronsneutronschargecharge
nucleusnucleus
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Chadwick’s Experiment (1932)(1935 Noble Prize in Physics)
H atoms: 1 p; He atoms: 2 p
mass He/mass H should = 2
measured mass He/mass H = 4
+ 9Be 1n + 12C + energy
neutron (n) is neutral (charge = 0)
n mass ~ p mass = 1.67 x 10-24 g
Parts of Atoms, Isotopes, and Ions
Fundamental parts of atoms
COMPONENTCHARGE
APPROXIMATE MASS (in amu)
LOCATION IN ATOM
electronelectron - 1- 1 0.000550.00055
shellshellss
protonproton + 1+ 1 1.0 1.0 nucleusnucleus
neutronneutron 00 1.000551.00055 nucleusnucleus
Parts of Atoms, Isotopes, + Ions
–Individual _____ of any element
–Represented with a _______ symbol
XXAA
ZZ
CC
atomatomnuclearnuclear
element symbolelement symbolmass numbermass number
atomic numberatomic number
•X = _______ _______
•A = ____ _______
•Z = ______ _______
Parts of Atoms, Isotopes, and Ions
•X = _______ _______
•A = ____ _______
•Z = ______ _______
= ______ _______
•C = ______ __ __ ___
•# neutrons = ___ - ___
•# electrons = ___ - ___
XXAA
ZZ
CCelement symbolelement symbolmass numbermass number
atomic numberatomic numberprotons numberprotons number
charge as an ioncharge as an ionAA ZZ
ZZ CC
Parts of Atoms, Isotopes, and Ions
XX1919 1-1-Mass number?Mass number?
Atomic number?Atomic number?
# of neutrons?# of neutrons?
Element?Element?
Charge?Charge?
# of electrons?# of electrons?
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Parts of Atoms, Isotopes, and Ions
• Isotopes
–Atoms of one element
•Same number of _______
•Different numbers of ________
•Different ______
protonsprotonsneutronsneutrons
massesmasses
Parts of Atoms, Isotopes, and Ions
• Ions
–________ atoms that have gained or lost ________
–Examples
• Iron loses 2 electrons
•Chlorine gains 1 electron
ChargedChargedelectronselectrons
FeFe +2 +2
ClCl -1 -1
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mass p ≈ mass n ≈ 1840 x mass e-
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Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei
XAZ
H11 H (D)2
1 H (T)31
U23592 U238
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Mass Number
Atomic NumberElement Symbol
Atomic Number, Mass Number, and Isotopes
57
The Isotopes of Hydrogen
Example 2.1
Give the number of protons, neutrons, and electrons in each of the following species:
(a)
(b)
(c)
(d) carbon-14
Example 2.1
Solution
(a) The atomic number is 11, so there are 11 protons. The mass number is 20, so the number of neutrons is 20 − 11 = 9. The number of electrons is the same as the number of protons; that is, 11.
(b) The atomic number is the same as that in (a), or 11. The mass number is 22, so the number of neutrons is 22 − 11 = 11. The number of electrons is 11. Note that the species in (a) and (b) are chemically similar isotopes of sodium.
Example 2.1
(c) The atomic number of O (oxygen) is 8, so there are 8 protons. The mass number is 17, so there are 17 − 8 = 9 neutrons. There are 8 electrons.
(d) Carbon-14 can also be represented as 14C. The atomic number of carbon is 6, so there are 14 − 6 = 8 neutrons. The number of electrons is 6.
Example
Atomic number = # of protons = # of electrons (neutral atom)
Atomic mass
Carbon has an atomic number of 6, which means it has 6 ________ in its nucleus and 6 __________ orbiting around the nucleus.
protonselectrons
Organization of Your Periodic Table
He, Ne, and Ar are in the same ___________.
C, N, and O are in the same ___________.
P E R I O D SGROUPS
metalloids
Metals
Excellent conductors (Sea of electrons) even as solids
Mobile electrons
Metals are malleable and ductile
Can be pounded into sheet
Can be pulled into wire
Alkali Metals
Similar properties:
Group # 1
Reactive
Good Conductors
Alkaline Earth Metals
Similar properties:
Group # 2
Reactive – but less than alkali
Good Conductors
Noble Gases
Similar properties:
Group # 18
NON-reactive
Gases
Full outer shells!!!
Halogens
Similar properties:
Group # 17
Very reactive
Fluorine and Chlorine = gases
7 valence electrons
Transition metalsSimilar properties:
Generally least reactive metals
Tricky valence (can change from the expected)
Mostly form +1, +2, and +3 ions
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The Modern Periodic Table
PeriodG
roup
Alkali M
etal
Noble G
as
Halogen
Alkali E
arth Metal
72
Chemistry In ActionNatural abundance of elements in Earth’s crust:
Natural abundance of elements in human body:
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A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical forces.
H2 H2O NH3 CH4
A diatomic molecule contains only two atoms:
H2, N2, O2, Br2, HCl, CO
A polyatomic molecule contains more than two atoms:
O3, H2O, NH3, CH4
diatomic elements
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An ion is an atom, or group of atoms, that has a net positive or negative charge.
cation – ion with a positive chargeIf a neutral atom loses one or more electronsit becomes a cation.
anion – ion with a negative chargeIf a neutral atom gains one or more electronsit becomes an anion.
Na 11 protons11 electrons Na+ 11 protons
10 electrons
Cl 17 protons17 electrons Cl-
17 protons18 electrons
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A monatomic ion contains only one atom:
A polyatomic ion contains more than one atom:
Na+, Cl-, Ca2+, O2-, Al3+, N3-
OH-, CN-, NH4+, NO3
-
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Common Ions Shown on the Periodic Table
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Formulas and Models
78
A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance.
An empirical formula shows the simplest whole-number ratio of the atoms in a substance.
H2OH2O
molecular empirical
C6H12O6 CH2O
O3 O
N2H4 NH2
Example 2.2
Write the molecular formula of methanol, an organic solvent and antifreeze, from its ball-and-stick model, shown below.
Example 2.2
Solution
Refer to the labels (also see back endpapers).
There are four H atoms, one C atom, and one O atom. Therefore, the molecular formula is CH4O.
However, the standard way of writing the molecular formula for methanol is CH3OH because it shows how the atoms are joined
in the molecule.
Example 2.3
Write the empirical formulas for the following molecules:
(a)acetylene (C2H2), which is used in welding torches
(b)glucose (C6H12O6), a substance known as blood sugar
(c)nitrous oxide (N2O), a gas that is used as an anesthetic gas (“laughing gas”) and as an aerosol propellant for whipped creams.
Example 2.3
Strategy
Recall that to write the empirical formula, the subscripts in the molecular formula must be converted to the smallest possible whole numbers.
Example 2.3
Solution
(a)There are two carbon atoms and two hydrogen atoms in acetylene. Dividing the subscripts by 2, we obtain the empirical formula CH.
(b) In glucose there are 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. Dividing the subscripts by 6, we obtain the empirical formula CH2O. Note that if we had divided the subscripts by 3, we would have obtained the formula C2H4O2.
Although the ratio of carbon to hydrogen to oxygen atoms in C2H4O2 is the same as that in C6H12O6 (1:2:1), C2H4O2 is not the
simplest formula because its subscripts are not in the smallest whole-number ratio.
Example 2.3
(c) Because the subscripts in N2O are already the smallest
possible whole numbers, the empirical formula for nitrous oxide is the same as its molecular formula.
85
Ionic compounds consist of a combination of cations and anions.
• The formula is usually the same as the empirical formula.
• The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero.
The ionic compound NaCl
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The most reactive metals (green) and the most reactive nonmetals (blue) combine to form ionic compounds.
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Formulas of Ionic Compounds
Al2O3
2 x +3 = +6 3 x -2 = -6
Al3+ O2-
CaBr2
1 x +2 = +2 2 x -1 = -2
Ca2+ Br-
Na2CO3
2 x +1 = +2 1 x -2 = -2
Na+ CO32-
Example 2.4
Write the formula of magnesium nitride, containing the Mg2+ and N3− ions.
When magnesium burns in air, it forms both magnesium oxide and magnesium nitride.
Example 2.4
Strategy Our guide for writing formulas for ionic compounds is electrical neutrality; that is, the total charge on the cation(s) must be equal to the total charge on the anion(s).
Because the charges on the Mg2+ and N3− ions are not equal, we know the formula cannot be MgN.
Instead, we write the formula as MgxNy, where x and y are
subscripts to be determined.
Example 2.4
Solution To satisfy electrical neutrality, the following relationship must hold:
(+2)x + (−3)y = 0
Solving, we obtain x/y = 3/2. Setting x = 3 and y = 2, we write
Check The subscripts are reduced to the smallest whole- number ratio of the atoms because the chemical formula of an ionic compound is usually its empirical formula.
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Chemical Nomenclature
• Ionic Compounds– Often a metal + nonmetal– Anion (nonmetal), add “-ide” to element name
BaCl2 barium chloride
K2O potassium oxide
Mg(OH)2 magnesium hydroxide
KNO3 potassium nitrate
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• Transition metal ionic compounds– indicate charge on metal with Roman numerals
FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride
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
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Example 2.5
Name the following compounds:
(a)Cu(NO3)2
(b)KH2PO4
(c)NH4ClO3
Example 2.5
Strategy Note that the compounds in (a) and (b) contain both metal and nonmetal atoms, so we expect them to be ionic compounds.
There are no metal atoms in (c) but there is an ammonium group, which bears a positive charge. So NH4ClO3 is also an
ionic compound.
Our reference for the names of cations and anions is Table 2.3.
Keep in mind that if a metal atom can form cations of different charges (see Figure 2.11), we need to use the Stock system.
Example 2.5
Solution
(a)The nitrate ion ( ) bears one negative charge, so the copper ion must have two positive charges. Because copper forms both Cu+ and Cu2+ ions, we need to use the Stock system and call the compound copper(II) nitrate.
(b)The cation is K+ and the anion is (dihydrogen phosphate). Because potassium only forms one type of ion (K+), there is no need to use potassium(I) in the name. The compound is potassium dihydrogen phosphate.
(c) The cation is (ammonium ion) and the anion is . The compound is ammonium chlorate.
Example 2.6
Write chemical formulas for the following compounds:
(a)mercury(I) nitrite
(b)cesium sulfide
(c)calcium phosphate
Example 2.6
Strategy
We refer to Table 2.3 for the formulas of cations and anions.
Recall that the Roman numerals in the Stock system provide useful information about the charges of the cation.
Example 2.6
Solution
(a)The Roman numeral shows that the mercury ion bears a +1 charge. According to Table 2.3, however, the mercury(I) ion is diatomic (that is, ) and the nitrite ion is . Therefore, the formula is Hg2(NO2)2.
(b)Each sulfide ion bears two negative charges, and each cesium ion bears one positive charge (cesium is in Group 1A, as is sodium). Therefore, the formula is Cs2S.
Example 2.6
(c) Each calcium ion (Ca2+) bears two positive charges, and each phosphate ion ( ) bears three negative charges.
To make the sum of the charges equal zero, we must adjust the numbers of cations and anions:
3(+2) + 2(−3) = 0
Thus, the formula is Ca3(PO4)2.
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• Molecular compounds− Nonmetals or nonmetals + metalloids− Common names
− H2O, NH3, CH4
− Element furthest to the left in a period and closest to the bottom of a group on periodic table is placed first in formula
− If more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom
− Last element name ends in -ide
103
HI hydrogen iodide
NF3 nitrogen trifluoride
SO2 sulfur dioxide
N2Cl4 dinitrogen tetrachloride
NO2 nitrogen dioxide
N2O dinitrogen monoxide
Molecular Compounds
Example 2.7
Name the following molecular compounds:
(a)SiCl4
(b)P4O10
Example 2.7
Strategy
We refer to Table 2.4 for prefixes.
In (a) there is only one Si atom so we do not use the prefix “mono.”
Solution
(a)Because there are four chlorine atoms present, the compound is silicon tetrachloride.
(b)There are four phosphorus atoms and ten oxygen atoms present, so the compound is tetraphosphorus decoxide. Note that the “a” is omitted in “deca.”
Example 2.8
Write chemical formulas for the following molecular compounds:
(a)carbon disulfide
(b) disilicon hexabromide
Example 2.8
Strategy
Here we need to convert prefixes to numbers of atoms (see Table 2.4).
Because there is no prefix for carbon in (a), it means that there is only one carbon atom present.
Solution
(a)Because there are two sulfur atoms and one carbon atom present, the formula is CS2.
(b) There are two silicon atoms and six bromine atoms present, so the formula is Si2Br6.
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109
An acid can be defined as a substance that yields hydrogen ions (H+) when dissolved in water.
For example: HCl gas and HCl in water
•Pure substance, hydrogen chloride
•Dissolved in water (H3O+ and Cl−), hydrochloric acid
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111
An oxoacid is an acid that contains hydrogen, oxygen, and another element.
HNO3 nitric acid
H2CO3 carbonic acid
H3PO4 phosphoric acid
112
Naming Oxoacids and Oxoanions
113
The rules for naming oxoanions, anions of oxoacids, are as follows:1. When all the H ions are removed from the
“-ic” acid, the anion’s name ends with “-ate.” 2. When all the H ions are removed from the
“-ous” acid, the anion’s name ends with “-ite.”3. The names of anions in which one or more
but not all the hydrogen ions have been removed must indicate the number of H ions present.
For example:– H2PO4
- dihydrogen phosphate– HPO4 2- hydrogen phosphate– PO4
3- phosphate
114
Example 2.9
Name the following oxoacid and oxoanion:
(a)H3PO3
(b)
Example 2.9
Strategy To name the acid in (a), we first identify the reference acid, whose name ends with “ic,” as shown in Figure 2.15.
In (b), we need to convert the anion to its parent acid shown in Table 2.6.
Solution
•We start with our reference acid, phosphoric acid (H3PO4). Because H3PO3 has one fewer O atom, it is called phosphorous acid.
•The parent acid is HIO4. Because the acid has one more O atom than our reference iodic acid (HIO3), it is called periodic acid. Therefore, the anion derived from HIO4 is called periodate.
117
A base can be defined as a substance that yields hydroxide ions (OH-) when dissolved in water.
NaOH sodium hydroxide
KOH potassium hydroxide
Ba(OH)2 barium hydroxide
118
Hydrates are compounds that have a specific number of water molecules attached to them.
BaCl2•2H2O
LiCl•H2O
MgSO4•7H2O
Sr(NO3)2 •4H2O
barium chloride dihydrate
lithium chloride monohydrate
magnesium sulfate heptahydrate
strontium nitrate tetrahydrate
CuSO4•5H2O CuSO4
119
120
Organic chemistry is the branch of chemistry that deals with carbon compounds.
C
H
H
H OH C
H
H
H NH2 C
H
H
H C OH
O
methanol methylamine acetic acid
Functional Groups:
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