chapter 4: atoms
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Chapter 4: Atoms. The Building Blocks of Matter. An atom is the smallest particle of an element that retains the chemical properties of that element. Section 1. The Atom: From Philosophical Idea to Scientific Theory Page 67. The Early Atom. - PowerPoint PPT PresentationTRANSCRIPT
Chapter 4: Atoms
The Building Blocks of Matter
An atom is the smallest particle of an element that retains the chemical properties of that
element.
Section 1
The Atom: From Philosophical Idea to Scientific Theory
Page 67
The Early Atom
As early as 400 B.C., Democritus called nature’s basic particle the “atomon” based on the Greek word meaning “indivisible”.
Aristotle succeeded Democritus and did not believe in atoms. Instead, he thought that all matter was continuous. It was his theory that was accepted for the next 2000 years. (Read page 43 of your textbook.)
Three Basic Laws of Matter:
Law of Conservation of MassLaw of Definite ProportionsLaw of Multiple Proportions
Basic Laws of Matter
Law of Conservation of Mass- mass is neither created nor destroyed during ordinary chemical reactions or physical changes.
CH4 + 2O2 → 2H2O + CO2
16g + 64g → 36g + 44g
Antoine Lavoisier
stated this about 1785
Antoine Lavoisier and his wife, Marie-Anne
"It took them only an instant to cut off that head, and a hundred years may not produce another like it." Joseph-Louis Lagrange
Alka Seltzer in Water Ziploc bag Alka seltzer tablet Water
Using the reaction between the tablet and the water, prove that the Law of Conservation of matter is true.
HOMEWORK
Read Section 1 Complete Questions 1-3 of the Section 1
Review (page 71) on a separate sheet of paper to be collected.
Basic Laws of Matter Law of Definite Proportions – no matter how much
salt you have, it is always 39.34% Na and 60.66% Cl by mass.
Example: Sodium chloride always contains 39.34% Na and 60.66% Cl by mass.
2NaCl → 2Na + Cl2100g → 39.34g + 60.66g116.88g → ? + ?
Joseph Louis Proust
stated this in 1794.
Basic Laws of Matter Law of Multiple Proportions- Two or more
elements can combine to form different compounds in whole-number ratios.
Example
John Dalton proposed this
in 1803.
John Dalton’s Elements
Dalton’s Atomic Theory In 1808, Dalton proposed a theory to
summarize and explain the laws of conservation of mass, definite proportions, & multiple proportions.
I was a school teacher at the
age of 12!
Dalton’s Atomic Theory
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.**
3. Atoms cannot be subdivided, created, or destroyed.**
4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.
**Today, we know these parts to have flaws.
John Dalton - 1808
Flaws of Dalton’s Theory…
2. Atoms of a given element are identical in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
Isotopes – atoms with the same number of protons but a different number of neutrons
Subatomic particles – electrons, protons, neutrons, and more
Section 2
The Structure of the Atom Page 72
The Atom
Atom - the smallest particle of an element that retains the chemical properties of that element.
CARBON
The Structure of the Atom
The atom is composed of two main regions, the nucleus & the electron cloud.
Nucleus of an Atom Nucleus- very small region located at the center of the
atom. The nucleus accounts for most of an atoms mass but very little volume, making it a very dense region.
The nucleus contains protons, neutrons, and more.
proton = p+ neutron = no others – neutral, too
MD=
V
Electron Cloud of an Atom The electron cloud is the negatively
charged region of the atom that accounts for most of the atom’s volume but very little of the atom’s mass.
electron = e-
The electron cloud is composed of a number of electrons, of which depends the element.
Checking for Understanding
What are the two main regions of the atom?
Does an electron from gold, act like gold?
What is the charge on the nucleus?
NO, an electron is like any other electron, no matter the source.
The nucleus and the electron cloud are the two main regions.
The nucleus is positive since it holds protons (+), neutrons (0) and other neutral particles.
Subatomic Particles
Protons- positively charged particles found in the nucleus of an atom.
Neutrons- neutral particles found in the nucleus of an atom.
Electrons- negatively charged particles found in the electron cloud.
Others – photon, boson, gluon, lepton, muon, quark, tau, neutrino, meson, …
Properties of Subatomic Particles
Particle Symbol Charge Mass # Relative Mass (amu)
Actual Mass (g)
Electron e- -1 0 0.0005486 9.109 X 10-28
Proton p+ +1 1 1.007276 1.673 X 10-24
Neutron no 0 1 1.008665 1.675 X 10-24
1 amu (atomic mass unit) = 1.660540 x 10-27 kg or exactly 1/12 the mass of a carbon-12 atom
Discovery of the Subatomic Particles
The discovery of the subatomic particles came about from the study of electricity & matter.
Benjamin Franklin’s kite experiment in 1752 demonstrated that lightning was electrical.
Charged Particles
In 1832, Michael Faraday proposed that objects are made of positive and negative charges.
Discovery of the Electron In the late 1870’s many experiments were
performed in which electric current was passed through gases at low pressures due to the fact that gases at atmospheric pressure don’t conduct electricity well.
These experiments were carried out in glass tubes called cathode-ray tubes or Crookes tubes.
Sir William Crookes developed these tubes.
Crookes TubeCRT
Discovery of the Electron
When current was passed through the cathode ray tube, the surface of the tube, directly opposite the cathode, glowed.
It was thought that this glow was caused by a stream of particles called cathode rays.
The rays traveled from cathode (negative) to anode (positive).
Discovery of the Electron
Negatively charged objects deflected the rays away.
Therefore, it was determined that the particles making up the cathode rays were negatively charged.
Joseph John Thomson
In 1897 the English physicist Joseph John Thomson was able to measure the ratio of charge of the cathode ray particles to their mass.
He found that the ratio was always the same regardless of the metal used to make the cathode or the nature of the gas inside the cathode ray tube.
Thomson concluded that cathode rays were composed of identical, negatively charged particles called electrons.
Joseph John Thomson
Thomson’s experiments revealed that the electron has a very large charge-to-mass ratio.
Thomson determined that electrons were present in all elements because he noted that cathode rays had identical properties regardless of the element used to produce them.
Cathode Ray Tube Experiment Accomplishments
Proved that the atom was divisible and that all atoms contain electrons.
This contradicted Dalton’s Atomic Theory. This allowed a new model of the atom.
Plum-Pudding Model of the Atom
Checking for UnderstandingCathode Ray Tube Why were the cathode rays deflected?
Why did they assume there was a positive portion to the atom?
How did this contradict Dalton’s model of the atom?
They were negatively charged, so they were repelled from the negative plate and attracted to the positive plate.
They knew the atom was neutral, so by default, there must be a positive portion if there are negative particles.
Dalton stated that atoms cannot be subdivided. Electrons are subatomic particles.
Robert A. Millikan
In 1909, Robert Millikan performed an Oil Drop Experiment & calculated the charge of the electron.
Oil Drop Experiment Millikan dropped negatively charged microscopic
oil particles into a chamber containing metallic plates and viewed them with a microscope.
By applying voltage to the metallic plates, Millikan created an electric field.
He was able to suspend the oil droplets by adjusting the electric field to the appropriate strength and direction to overcome gravity.
Oil Drop Experiment
Knowing the mass of the droplets and the strength of the electric field necessary to suspend them, he was able to calculate the charge of the electron.
He noticed that the charge was always a whole-number multiple of 1.602 X10-19 Coulombs.
He determined that the charge of the electron to be 1.602 X 10-19 C.
Checking for UnderstandingOil Drop Experiment What year did Millikan perform this
experiment?
How did he view the oil droplets?
He did NOT measure the charge on the electron; he calculated it. What did he measure?
1909
He viewed them with a microscope.
He knew the mass of the droplets and the strength of the electric field.
Discovery of X-Rays
In 1895 William Conrad Roentgen discovered X-rays, a form of radiation.
Radioactivity In 1896, the French scientist
Henri Becquerel was studying a Uranium mineral. He discovered it was spontaneously emitting high-energy radiation.
In 1898, Marie and Pierre Curie attempted to isolate radioactive components of the mineral.
Radioactivity
In 1899, Ernest Rutherford, a British scientist, began to classify radiation: alpha (), beta (), and gamma ().
Radiation Look closely at the paths of radiation. Do
you notice something about the amount of deflection of each type of particles?
Radiation
Discovery of the Nucleus
In 1911, Ernest Rutherford performed a Gold Foil Experiment.
He and his colleagues bombarded a thin piece of gold foil with fast moving, positively charged alpha particles.
Alpha Particles
Alpha () particles are Helium-4 nuclei. This means they are two protons and two
neutrons (with no electrons). Thus, they are positive.
4 +22 He
Gold Foil Experiment
Gold Foil Experiment As expected, most of the alpha particles
passed straight through with little or no deflection.
However, 1/8000 of the positively charged alpha particles were deflected, some back at the source.
(Po)
Gold Foil Experiment
Gold Foil Experiment
From this experiment, Rutherford discovered that there must be a very densely packed positively charged bundle of matter within the atom which caused the deflections.
He called this positive bundle the nucleus. He tried this experiment with other metals
and found the same results.
Gold Foil Experiment The volume of the
nucleus was very small compared to the volume of the atom.
Therefore, most of the atom was composed of empty space. Niels Bohr later found that this empty space was where the electrons were located.
Checking for UnderstandingGold Foil Experiment Why did some of the alpha particles come
straight back to the source or deflect away from the nucleus?
Why did he conclude that the nucleus must be positive?
What things did Rutherford conclude from the gold foil experiment?
Checking for UnderstandingGold Foil Experiment If gold atoms were solid spheres stacked together
with no space between them, what would you expect would happen to particles shot at them?
What year did Ernest Rutherford perform this experiment?
Rutherford experimented with many kinds of metal foil as the target. The results were always similar. Why was it important to do this?
“It was about as believable as if you had fired a 15-inch
shell at a piece of tissue paper, and it came back
and hit you.”
-Ernest Rutherford
Discovery of the Neutron In 1932, James Chadwick discovered the
neutron. Rutherford predicted that there were
massive, neutrally charged particles in the nucleus, but it was Chadwick who proved their existence.
Bohr’s Model of The Atom
Forces in the Nucleus
The nucleus is positive (p+ and no) Like charges repel each other…so
shouldn’t the p+ in the nucleus repel each other?
But…when 2 p+ are close together in the nucleus there is a strong attraction between them. The same holds true for neutrons.
REMEMBER:
Forces in the Nucleus
no act like the “glue” that holds the nucleus together. They help to stabilize the nucleus.
Nuclear forces are the short-range p+-no, p+- p+ , and no-no forces that hold the nuclear particles together.
Atomic Number
atomic number (Z) - the number of protons in the nucleus of each atom of a given element.
The number of p+ identifies the element.
Atomic Number increases from left to right on the periodic table.
Electrons
The number of electrons in a neutral atom is equal to the number of protons in that atom.
e- = p+
•Electrons can be lost or gained.
• When electrons are lost or gained, ions are formed.
Ionsion- an atom with a positive or negative
charge.cation- an atom with a positive charge Cations are formed when an atom loses
negatively charged electrons. Ca+2 is formed when calcium loses 2
electrons. Ca+2 has 2 less electrons than protons.
the lithium atom
the lithium ion
Ions
anion- an atom with a negative charge Anions are formed when an atom gains
negatively charged electrons. N-3 is formed when nitrogen gains 3
electrons. N-3 has 3 more electrons than protons.
Noble gases are very stable
and don’t react.
Every element on the periodic
table will try to react to be
stable, like the noble gases.
Metals vs. Nonmetals
Metals form cations.
Na Na+ + 1e-
Nonmetals form anions.
Cl + 1 e- Cl-
Charge determination
Group 1 – forms +1 Group 2 – forms +2 Group 13(B and Al) – forms +3 Group 15 – forms -3 Group 16 – forms -2 Group 17 – forms -1 Group 18 – doesn’t forms ions easily!
(WITH SOME EXCEPTIONS!)
Mass Number
mass number (A)- the number of p+ & no in the nucleus of an atom.
# of neutrons = mass number – atomic number
Why aren’t electrons included when determining the mass number of the atom?
Isotopes
isotope- two or more atoms having the same atomic number (same #p+) , but different mass numbers (due to different #no).
nuclide- general term for a specific isotope of an element.
Isotopes
Isotope Notation
Nuclear Notation
Hyphen NotationUses the elements symbol followed by a hyphen & the mass
number.
C-12
How many protons, neutrons & electrons are there in the following?
Cl-38 35Cl-1
Br-80 32S-2
N-14 56Fe+3
You Try It.
Do the Subatomic Particles Table Worksheet.
Changes in the Nucleus
Nuclear Reaction- changes that occur in the atom’s nucleus.
Nuclear reactions can change the composition of an atom’s nucleus permanently.
Types of Radiation Produced in Nuclear Reactions
Alpha ()
Beta ()
Gamma ()
Nuclear Stability
Atoms with unstable nuclei are radioactive. Most atoms have stable nuclei and are,
therefore, not radioactive.
Nuclear Stability
Neutrons help to stabilize the nucleus.
Elements 1-20 have p+ = no
Above element 20, increasingly more no are needed than p+ to maintain nuclear stability.
Element 84 and up, all atoms are radioactive so the nucleus cannot be stabilized regardless of the number of no.
Types of Radioactive Decay
Alpha Radiation ()- stream of high energy alpha particles. Consists of 2 protons & 2 neutrons making it identical to
a He-4 nucleus. Alpha particles can be represented by:
Most alpha particles are able to travel only a few centimeters through air and are easily stopped by clothing etc.
4 4 +2 42 2 2 He He
Alpha Decay
239 235 494 92 2Pu U + He
234 230 492 90 2U Th + He
parent daughter
Types of Radioactive Decay
Beta Radiation () – consists of a stream of high speed electrons. These electrons are not electrons that are in motion around the atom’s nucleus.
Beta particles can be represented by:
Can penetrate through clothing and damage skin.
0 -1 0 01 1 1 e e
Beta Decay
6 6 02 3 -1He Li + 24 24 011 12 -1Na Mg +
parent daughter
Types of Radioactive Decay
Gamma Rays ()- energetic form of light that cannot be seen.
Does not contain particles. Gamma particles can be represented by:
Can penetrate heavy material including skin. Can only be stopped by lead or concrete.
00
Checking for Understanding
21084 Po 14
6C
238 234 234 234 230U Th Pa U Th ?
alpha decay
beta decay
4 2062 82He + Pb
14 07 -1N +
22688 Ra
Other Types of Nuclear Reactions positron –
proton -
neutron -
01e
10 n
11H
Half Lifehalf life- the time required for half of the atoms in any given quantity of a radioactive isotope to decayEach particular isotope has its own half-life.
Half Life
p.689 Sample Problem B
Phosphorus-32 has a half-life of 14.3 days. How many milligrams of phosphorus-32 remain after 57.2 days if you start with 4.0 mg of the isotope?
Ans: 0.25 mg
You Try It.
Do the Nuclear Reactions Worksheet.
The Mole
mole (mol)- SI Unit for the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12.
A unit of counting, like the dozen.
Avogadro’s Number
Avogadro’s Number - the number of particles in exactly one mole of a pure substance.
1 mole = 6.0221415 X 1023
1 mol = 6.02 x 1023
Amedeo Avogadro
Atomic Mass
atomic mass - the mass of one mole of an atom
Atomic mass is expressed in atomic mass units (amu) or (u) or g/mol.
Can be found on the periodic table. All atomic masses are based on the
atomic mass of carbon-12 being 12 amu.
Molar Mass
molar mass - the mass of one mole of a pure substance.
Molar mass is written in units of amu or g/mol.
Atomic mass vs. Molar mass
atomic mass - the mass of one mole of an atom.
molar mass - the mass of one mole of a pure substance.
Atomic Mass vs. Molar Mass
Example Atomic Mass
Na
Ag
C
O
22.99 g/mol
107.87 g/mol
12.01 g/mol
16.00 g/mol
Molar Mass of Compounds
Compound Molar Mass
H2O
C6H12O6
NaCl
Cl2
(NH4)3PO4
CuSO4·5H2O
18.02 g/mol
180.18 g/mol
58.44 g/mol
70.90 g/mol
149.12 g/mol
249.72 g/mol
Introduction to Molar Conversions
Amount Mass
1 mol O2
½ mol O2
2 mol O2
3 mol O2
32.00 g
16.00 g
64.00 g
96.00 g
Cheer
I say grams, you say molar mass.
grams – molar mass
grams – molar mass
Grams to MolesConverting grams to moles: divide by molar mass.
1. How many moles of Ca are in 5.00g of Ca?
2. How many moles of H2O are in 36.0g of H2O?
3. How many moles of AgNO3 are in 124.5g of AgNO3?
1 mol Ca5.00g Ca x =
40.08 g Ca0.125 mol Ca
22
2
1 mol H O36.0 g H O x =
18.02 g H O22.00 mol H O
33
3
1 mol AgNO124.5 g AgNO x =
169.88 g AgNO0.7329
3mol AgNO
Moles to Grams
Converting from moles to grams: multiply by molar mass
1. What is the mass in grams of 2.25 moles of Fe?
2. What is the mass in grams of 0.896 moles of BaCl2?
55.85 g Fe2.25 mol Fe x =
1 mole Fe126 g Fe
2187 g BaCl22
2
208.23 g BaCl0.896 mol BaCl x =
1 mole BaCl
Types of Particles
Atoms – C, Cu, He
Molecules – O2, C12H22O11, CO2 (all nonmetals in the formula)
Formula units – NaCl, CaCl2, Mg(NO3)2 (includes a metal in the formula)
1 mole = 6.02 x 1023 particles
Particles to MolesConverting particles to moles: divide by
Avogadro's Number.
1. How many moles of Pb are in 1.50 X 1025 atoms of Pb?
2. How many moles of CO2 are in 6.78 X 1021 molecules of CO2?
2.49 x 101 moles Pb
1.13 x 10-2 moles CO2
Moles to ParticlesConverting moles to atoms: multiply by
Avogadro's Number.
1. How many molecules of NO are in 0.87 moles of NO?
2. How many formula units of NaI are in 2.50 moles of NaI?
5.2 x 1023 molecules NO
1.51 x 1024 formula units NaI
Grams to Moles to Particles
Example: How many molecules of N2 are in 57.1g of N2?
257.1 g N x 2
2
1 mol N
28.02 g N
232
2
6.02 x 10 molecules Nx
1 mol N
242= 1.23 x 10 molecules N
Particles to Moles to Grams
Example: How many grams of NaF are in 7.89 X 1024 formula units of NaF?
247.89 x 10 f.un. NaF x 23
1 mol NaF
6.02 x 10 f.un. NaF
41.99 g NaFx
1 mol NaF
= 550. g NaF
Atoms to Moles to Grams
Tough Example:
How many total atoms are in 235 g of CO2?
9.64 x 1024 total atoms
The Mole Bridge
Atomic Mass Determination
Average Atomic Mass - the weighted average of atomic masses of the naturally occurring isotopes of an element.
The atomic mass is expressed relative to the value of exactly 12u for a carbon-12 atom.
Atomic Mass Unit – amu or u
What is a weighted average?
Example: Your grade in math might be 75% tests and 25% homework. What would your grade be if you had a test average of 80% and a homework average of 100%?
Normally, you would average 80% & 100% to get 90%.
However, with a weighted average, it is 85%.
80%(.75) 100%(.25) 85%
Calculating Average Atomic Mass
Average atomic mass =
atomic mass of each isotope X percent natural abundance (in decimal form)
average atomic mass =
(atomic mass of each isotope x % abundance of each isotope in decimal form)
Percent Natural Abundance
Percent Natural Abundance- the relative proportions expressed as percentages, in which isotopes of an element are found in nature.
Calculating Average Atomic Mass
Ex: Uranium has two isotopes, uranium-235 and uranium-238. Uranium-235 has an atomic mass of 235.043amu and a percent natural abundance of 0.720%. Uranium-238 has an atomic mass of 238.050amu and a percent natural abundance of 99.280%. Calculate the average atomic mass of naturally occurring uranium.