Chapter 2. Atoms, Molecules, and Ions
Chapter 2
Atoms, Molecules, and Ions
DE Chemistry
Dr. Walker
Timeline
B.C.
400 B.C. Demokritos and Leucippos
use the term "atomos”
1500's
Georg Bauer: systematic metallurgy
Paracelsus: medicinal application of minerals
1600's
Robert Boyle:
The Skeptical Chemist.
Quantitative experimentation, identification of
elements
1700s'
Georg Stahl: Phlogiston Theory
Joseph Priestly: Discovery of oxygen
Antoine Lavoisier: The role of oxygen in combustion, law of conservation of
mass, first modern chemistry textbook
2000 years of Alchemy (other elements to gold)
Timeline
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Dalton’s Atomic Theory (1808) – The
Dalton’s Atomic Theory (1808) – The
Short Version
Short Version
•
Each element is made up of tiny particles
called atoms
•
The atoms of a given element are
identical
•
Chemical compounds are formed when
atoms combine with each other. A given
compound always has the same relative
numbers and types of atoms
•
Chemical reactions involve reorganizations
of the atoms. The atoms themselves are
not changed in a chemical reaction
Avogadro (Yes, him again…)
•
At the same conditions of temperature and
pressure, equal volumes of different gases
contain the same number of particles.
–
Relates to the mole concept
•
1 mole = 22.4 L (@ STP) = molar mass (grams) = 6.02 x
10
23
particles
–
Relates to gas laws
•
1 mole
not equal to 22.4 L at other conditions
•
Find volume by using ideal gas law
Thomson (1897)
•
Cathode Ray Tube Experiment
–
Determined the charge to mass ratio of the
electron
–
Reasoned that all atoms must contain electrons
–
Reasoned that all atoms must contain positive
charges
Thomson’s Atomic Model
•
Thomson believed that the
electrons were like plums
embedded in a positively
charged “pudding,” thus it was
called the “plum pudding”
model.
J.J. Thomson
J.J. Thomson
Millikan – Mass of Electrons
•
1909 – Robert Millikan determines the
mass of the electron using charged oil
droplets
–
Mass of the electron is 9.109 x 10
-31
kg
Fundamental Chemical Laws
•
Law of Conservation of Mass
–
"Mass is neither created nor destroyed"
–
Translation: In ordinary chemical reactions, the
total mass of the reactants is equal to the total
mass of the products
•
Why equations need to be balanced
•
Basis for stoichiometry, percent yield, etc.
Fundamental Chemical Laws
•
Law of Definite Proportion
–
"A given compound always contains the same
proportions of elements by mass"
–
Translation: Compounds have an unchanging
chemical formula
•
Different Formula, different compound
Fundamental Chemical Laws
•
Law of Multiple Proportions (Dalton)
–
When two elements form a series of compounds, the
ratios of the masses of the second element that
combine with one gram of the first element can
always be reduced to small whole numbers
–
Translation: Sometimes two elements can come
together in more than one way, forming compounds
with similar, though not identical formulas
•
Example: Hydrogen and Oxygen
–
2 H
2
+ O
2
2 H
2
O
–
H
2
+ O
2
H
2
O
2
Radioactivity
•
Gamma (
) rays - high energy light
•
Beta (
) particles - high speed electrons
•
Alpha (
) particles – helium nucleus, particle
with a 2+ charge
Rutherford
•
Gold Foil Experiment
–
Alpha particles are helium nuclei
–
Particles were fired at a thin sheet of gold foil
–
Particle hits on the detecting screen (film) are
recorded
Gold Foil Experiment
•
Some particles were greatly deflected ("like ahowitzer shell bouncing off of a piece of
paper")
–
Could not have been deflected by electrons or
single protons
–
Must have been deflected by a positively charged
object of substantial mass
•
Supported concept of a small, central, positive nucleus
where most of the atom's mass was concentrated
•
Disproved Thomson's "plum pudding" model
Atomic Scale
•
Nucleus is roughly 1/10
5
the
diameter of the atom
Most of the mass of the atom is
in the nucleus (protons and
neutrons)
Electrons are found outside of
the nucleus (the electron cloud)
Most of the volume of the
atom is empty space
Nucleus = 10
-15
m
Atom = 10
-10
m
Quarks – The True Fundamental
Particle
•
Protons and neutrons are
NOT
fundamental particles.
•
Protons are made of two “up” quarks
and one “down” quark.
•
Neutrons are made of one “up”
quark and two “down” quarks.
•
Quarks are held together by “gluons”
Atomic Particles
Atomic Number
•
Atomic Number = # protons = # electrons
(when neutral)
•
Mass Number = protons + neutrons
Atomic Number
•
Atomic Number = # protons = # electrons
(when neutral)
•
Mass Number = protons + neutrons
Atomic Masses
•
Atomic mass is the
weighted average
weighted average
of all
the naturally isotopes of that element.
Carbon = 12.011
Isotopes
•
Isotopes – Same element
–
Same # protons, varying # neutrons
From Atoms to Compounds
DE Version
Molecules and Ions
•
Chemical Bonding
–
Covalent Bonding – Sharing of electrons
–
Ionic Bonding – Transfer of electrons between
elements which form ions
Covalently Bound Molecules
•
Compounds with
covalent bonds
referred to as
molecules
•
Formulas
represent each
atom present in
the molecule
Structural formula
Space filling model
Ball and Stick model
Ions
Cation: A positive ion formed by losing electrons
Mg
2+
, NH
4
+
Anion: A negative ion formed by gaining electrons
Cl
, SO
4
2
Remember, atoms want a full valence shell !
Ionic Bonding: Force of attraction between oppositely
charged ions.
Ionic compounds form
crystals
, so their formulas are
written empirically (lowest whole number ratio of
ions).
Valence Electrons and Charge
Keep in mind:
-This is a generalization. There are numerous exceptions.
-
Remember, many (not all) transition metals can have multiple
charges.
Common Ions
•
This is not a complete list, but includes general
examples
Binary Ionic Compounds
•
Binary = 2 elements
–
Cation first, use name of the element
–
Anion second, use root of the element name and
add
–ide
•
Na
+
and Cl
-
makes NaCl = sodium chlor
ide
•
Ca
2+
and Br
-
makes CaBr
2
= calcium brom
ide
•
K
+
and O
2-
makes K
2
O = potassium ox
ide
–
As before, you will be expected to go in both
directions
•
Name to formula
AND
formula to name!
Common Polyatomic Ions
Rules of Thumb – Naming Polyatomics
•
Find the "ate" ion (sulfate, for instance)
–
sulfate = SO
4
2-
•
The "ite" ion always has one less oxygen than the
"ate" ion
–
sulfite = SO
3
2-
•
The prefix "per" (think hyper, meaning "above") is
used with the "ate" prefix to indicate one more
oxygen than the "ate" ion
–
persulfate = SO
5
2-
•
The prefix "hypo" (meaning "under" or "below")
is used with the "ite" prefix to indicate one less
–
hyposulfite = SO
2
2-
Rules of Thumb – Naming Polyatomics
Examples
•
Chlorate and Nitrate have 3 oxygens
–
“Per”-”ate” has 4 oxygens (1 more)
–
“ite” has 2 oxygens (1 less)
–
“hypo” – “ite” has 1 oxygen (1 less than “ite”)
Practice
Practice - Answers
•
a. KI
•
b. CaO
•
C. GaBr
3
Binary Ionic Compounds – Transition
Metals*
•
Same rules apply with one exception
–
Use roman numeral to show positive charge on
cation
–
Fe
3+
and Cl
-
makes FeCl
3
= iron (III) chloride
–
Cu
+
and F
-
makes CuF = copper (II) fluoride
–
Pb
4+
and O
2-
makes PbO
2
= lead (IV) oxide*
•
Remember, reduce numbers in formulas when possible
•
*Metals such as tin and lead are not transition metals,
yet they can have multiple charges
Ionic Compounds - Polyatomics
•
Cation first
–
List cation, whether it is an element or polyatomic
•
Anion second (as usual)
–
Use
–ide
for elements, list name if polyatomic
•
Use parentheses for multiple polyatomics, as
shown below
•
NH
4+
and S
2-
makes (NH
4
)
2
S, ammonium sulfide
•
Ca
2+
and C
2
H
3
O
2
-
makes Ca(C
2
H
3
O
2
)
2
, calcium
acetate
Naming Acids
•
Binary Acids (two elements - hydrogen + one
other)
–
prefix "Hydro" + root of second element + "ic"
suffix
•
HCl = hydrochloric acid
•
HBr = hydrobromic acid
Naming Acids
•
Oxyacids – acid including oxygen (from
polyatomic)
–
If the acid contains an anion whose name ends in
"ate":
–
Use root of anion name and an "ic" ending
•
H
2
SO
4
= sulfuric acid (SO
4
= sulfate)
•
H
2
CO
3
= carbonic acid (CO
3
= carbonate)
–
If the acid contains an anion whose name ends in
"ite": Use the root of the anion name and an "ous"
ending
•
H
2
SO
3
= sulfurous acid (SO
3
= sulfite)
•
HClO = hypochlorous acid (ClO = hypochlorite)
Covalent Bonding
•
Naming covalent compounds:
–
Use prefixes
Molecular Compounds
Compounds between two nonmetals
First element in the formula is named first.
Keeps its element name
Gets a prefix
if
there is a subscript on it
Second element is named second
Use the root of the element name plus the
-ide
suffix
Always
use a prefix on the second element
–
SO
2
»
Sulfur dioxide
–
N
2
O
3
»
Dinitrogen trioxide
–
SiBr
4
»
Silicon tetrabromide
Practice
Answers
•
a. SF
2
•
B. SF
6
•
C. NaH
2
PO
4
•
D. Li
3
N
•
E. Cr
2
(CO
3
)
3
•
F. SnF
2
•
G. NH
4
C
2
H
3
O
2
•
H. NH
4
HSO
4
•
i. Co(NO
3
)
3
•
J. HgCl
•
K. KClO
3
•
L. NaH
Practice
Answers
a. Cesium fluoride c. Silver sulfide* e. Titanium (IV) oxide
b. Lithium Nitride d. Manganese (IV) oxide f. Strontium phosphide
* Despite being a transition metal, silver only forms one cation
a.
Barium sulfate c. Potassium permanganate
b.
Sodium nitrite d. Potassium dichromate
Problem Set
•
Chapter 2
•
# 7, 11, 12, 22, 24, 42, 44, 46, 50, 54, 56, 60,
64, 66
Explore the historical timeline of atomic theory advancements, from Demokritos' initial concept of "atomos" to Dalton's Atomic Theory, Thomson's discovery of the electron, and modern breakthroughs by scientists like Rutherford and Lawrence. Avogadro's contributions to the mole concept and gas laws are also highlighted, showcasing the progression of understanding atoms, molecules, and ions in chemistry.
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Presentation Transcript
Chapter 2 Atoms, Molecules, and Ions DE Chemistry Dr. Walker
Timeline B.C. 400 B.C. Demokritos and Leucippos use the term "atomos 2000 years of Alchemy (other elements to gold) 1500's Georg Bauer: systematic metallurgy Paracelsus: medicinal application of minerals 1600's Robert Boyle: The Skeptical Chemist. Quantitative experimentation, identification of elements 1700s' Georg Stahl: Phlogiston Theory Joseph Priestly: Discovery of oxygen Antoine Lavoisier: The role of oxygen in combustion, law of conservation of mass, first modern chemistry textbook
Timeline 1800's Joseph Proust: The law of definite proportion (composition) John Dalton: The Atomic Theory, The law of multiple proportions Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules Amadeo Avogadro: Molar volumes of gases Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements Dmitri Mendeleyev: The periodic table J.J. Thomson: discovery of the electron Henri Becquerel: Discovery of radioactivity 1900's Robert Millikan: Charge and mass of the electron Ernest Rutherford: Existence of the nucleus, and its relative size Meitner & Fermi: Sustained nuclear fission Ernest Lawrence: The cyclotron and trans-uranium elements
Daltons Atomic Theory (1808) The Short Version Each element is made up of tiny particles called atoms The atoms of a given element are identical Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms Chemical reactions involve reorganizations of the atoms. The atoms themselves are not changed in a chemical reaction
Avogadro (Yes, him again) At the same conditions of temperature and pressure, equal volumes of different gases contain the same number of particles. Relates to the mole concept 1 mole = 22.4 L (@ STP) = molar mass (grams) = 6.02 x 1023 particles Relates to gas laws 1 mole not equal to 22.4 L at other conditions Find volume by using ideal gas law
Thomson (1897) Cathode Ray Tube Experiment Determined the charge to mass ratio of the electron Reasoned that all atoms must contain electrons Reasoned that all atoms must contain positive charges
Thomsons Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged pudding, thus it was called the plum pudding model. J.J. Thomson
Millikan Mass of Electrons 1909 Robert Millikan determines the mass of the electron using charged oil droplets Mass of the electron is 9.109 x 10-31 kg
Fundamental Chemical Laws Law of Conservation of Mass "Mass is neither created nor destroyed" Translation: In ordinary chemical reactions, the total mass of the reactants is equal to the total mass of the products Why equations need to be balanced Basis for stoichiometry, percent yield, etc.
Fundamental Chemical Laws Law of Definite Proportion "A given compound always contains the same proportions of elements by mass" Translation: Compounds have an unchanging chemical formula Different Formula, different compound
Fundamental Chemical Laws Law of Multiple Proportions (Dalton) When two elements form a series of compounds, the ratios of the masses of the second element that combine with one gram of the first element can always be reduced to small whole numbers Translation: Sometimes two elements can come together in more than one way, forming compounds with similar, though not identical formulas Example: Hydrogen and Oxygen 2 H2 + O2 2 H2O H2 + O2 H2O2
Radioactivity Gamma ( ) rays - high energy light Beta ( ) particles - high speed electrons Alpha ( ) particles helium nucleus, particle with a 2+ charge
Rutherford Gold Foil Experiment Alpha particles are helium nuclei Particles were fired at a thin sheet of gold foil Particle hits on the detecting screen (film) are recorded
Gold Foil Experiment Some particles were greatly deflected ("like a howitzer shell bouncing off of a piece of paper") Could not have been deflected by electrons or single protons Must have been deflected by a positively charged object of substantial mass Supported concept of a small, central, positive nucleus where most of the atom's mass was concentrated Disproved Thomson's "plum pudding" model
Atomic Scale Nucleus is roughly 1/105 the diameter of the atom Most of the mass of the atom is in the nucleus (protons and neutrons) Electrons are found outside of the nucleus (the electron cloud) Most of the volume of the atom is empty space Nucleus = 10-15 m Atom = 10-10 m
Quarks The True Fundamental Particle Protons and neutrons are NOT fundamental particles. Protons are made of two up quarks and one down quark. Neutrons are made of one up quark and two down quarks. Quarks are held together by gluons
Atomic Particles Subatomic Particle Mass Location Charge Proton 1.67 x 10-27 kg Nucleus + 1 Neutron 1.67 x 10-27 kg Nucleus 0 Electron 9.1 x 10-31 kg (much smaller) Orbitals (outside nucleus) - 1
Atomic Number Atomic Number = # protons = # electrons (when neutral) Mass Number = protons + neutrons Nuclide p+ n0 e- Mass # Oxygen - 10 - 33 42 - 31 15
Atomic Number Atomic Number = # protons = # electrons (when neutral) Mass Number = protons + neutrons Nuclide p+ n0 e- Mass # Oxygen - 18 8 10 8 18 Arsenic - 75 33 42 33 75 Phosphorous - 31 15 16 15 31
Atomic Masses Atomic mass is the weighted average of all the naturally isotopes of that element. Isotope Symbol Composition of % in nature the nucleus 6 protons 6 neutrons Carbon-12 12C 98.89% Carbon-13 13C 6 protons 7 neutrons 1.11% Carbon-14 14C 6 protons 8 neutrons <0.01% Carbon = 12.011
Isotopes Isotopes Same element Same # protons, varying # neutrons Isotope Protons Electrons Neutrons Nucleus Hydrogen 1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 1 1 Hydrogen-3 (tritium) 1 1 2
From Atoms to Compounds DE Version
Molecules and Ions Chemical Bonding Covalent Bonding Sharing of electrons Ionic Bonding Transfer of electrons between elements which form ions
Covalently Bound Molecules Compounds with covalent bonds referred to as molecules Formulas represent each atom present in the molecule Structural formula Space filling model Ball and Stick model
Ions Cation: A positive ion formed by losing electrons Mg2+, NH4+ Anion: A negative ion formed by gaining electrons Cl , SO42 Remember, atoms want a full valence shell ! Ionic Bonding: Force of attraction between oppositely charged ions. Ionic compounds form crystals, so their formulas are written empirically (lowest whole number ratio of ions).
Valence Electrons and Charge Group 1 2 13 14 15 16 17 18 Valence Electrons 1 2 3 4 5 6 7 8 Lose or Gain Lose 1 Lose 2 Lose 3 Either/ or Gain 3 Gain 2 Gain 1 Neither Charge +1 +2 +3 +4 or - 4 -3 -2 -1 0 Keep in mind: -This is a generalization. There are numerous exceptions. - Remember, many (not all) transition metals can have multiple charges.
Common Ions This is not a complete list, but includes general examples
Binary Ionic Compounds Binary = 2 elements Cation first, use name of the element Anion second, use root of the element name and add ide Na+ and Cl- makes NaCl = sodium chloride Ca2+ and Br- makes CaBr2 = calcium bromide K+ and O2- makes K2O = potassium oxide As before, you will be expected to go in both directions Name to formula AND formula to name!
Rules of Thumb Naming Polyatomics Find the "ate" ion (sulfate, for instance) sulfate = SO42- The "ite" ion always has one less oxygen than the "ate" ion sulfite = SO32- The prefix "per" (think hyper, meaning "above") is used with the "ate" prefix to indicate one more oxygen than the "ate" ion persulfate = SO52- The prefix "hypo" (meaning "under" or "below") is used with the "ite" prefix to indicate one less hyposulfite = SO22-
Rules of Thumb Naming Polyatomics Examples Chlorate and Nitrate have 3 oxygens Per - ate has 4 oxygens (1 more) ite has 2 oxygens (1 less) hypo ite has 1 oxygen (1 less than ite )
Practice - Answers a. KI b. CaO C. GaBr3
Binary Ionic Compounds Transition Metals* Same rules apply with one exception Use roman numeral to show positive charge on cation Fe3+ and Cl- makes FeCl3 = iron (III) chloride Cu+ and F- makes CuF = copper (II) fluoride Pb4+ and O2- makes PbO2 = lead (IV) oxide* Remember, reduce numbers in formulas when possible *Metals such as tin and lead are not transition metals, yet they can have multiple charges
Ionic Compounds - Polyatomics Cation first List cation, whether it is an element or polyatomic Anion second (as usual) Use ide for elements, list name if polyatomic Use parentheses for multiple polyatomics, as shown below NH4+ and S2- makes (NH4)2S, ammonium sulfide Ca2+ and C2H3O2- makes Ca(C2H3O2)2, calcium acetate
Naming Acids Binary Acids (two elements - hydrogen + one other) prefix "Hydro" + root of second element + "ic" suffix HCl = hydrochloric acid HBr = hydrobromic acid
Naming Acids Oxyacids acid including oxygen (from polyatomic) If the acid contains an anion whose name ends in "ate": Use root of anion name and an "ic" ending H2SO4 = sulfuric acid (SO4 = sulfate) H2CO3 = carbonic acid (CO3 = carbonate) If the acid contains an anion whose name ends in "ite": Use the root of the anion name and an "ous" ending H2SO3 = sulfurous acid (SO3 = sulfite) HClO = hypochlorous acid (ClO = hypochlorite)
Covalent Bonding Naming covalent compounds: Use prefixes
Molecular Compounds Compounds between two nonmetals First element in the formula is named first. Keeps its element name Gets a prefix if there is a subscript on it Second element is named second Use the root of the element name plus the -ide suffix Always use a prefix on the second element SO2 Sulfur dioxide N2O3 Dinitrogen trioxide SiBr4 Silicon tetrabromide
Answers a. SF2 B. SF6 C. NaH2PO4 D. Li3N E. Cr2(CO3)3 F. SnF2 G. NH4C2H3O2 H. NH4HSO4 i. Co(NO3)3 J. HgCl K. KClO3 L. NaH
Answers a. Cesium fluoride c. Silver sulfide* e. Titanium (IV) oxide b. Lithium Nitride d. Manganese (IV) oxide f. Strontium phosphide * Despite being a transition metal, silver only forms one cation a. Barium sulfate c. Potassium permanganate b. Sodium nitrite d. Potassium dichromate
Problem Set Chapter 2 # 7, 11, 12, 22, 24, 42, 44, 46, 50, 54, 56, 60, 64, 66
