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Chapter 2
Atoms and
Elements
AP CHEMISTRY
Copyright  2011 Pearson Education, Inc.
Dalton’s Atomic Theory
1. Each element is composed of tiny, indestructible
2.
3.
4.
particles called atoms
All atoms of an element have the same mass and
properties that distinguish them from atoms of
other elements
Atoms combine in simple, whole-number ratios to
form molecules
In a chemical reaction, atoms of one element
cannot change into atoms of another element, they
simply rearrange the way they are attached to
each other
2
Decide which statement is correct according
to Dalton’s model of the atom
1. Copper atoms can combine with
2.
3.
4.
3
zinc atoms to make gold atoms
Bulk water is composed of many
identical molecules, each having
one oxygen atom and two hydrogen
atoms
Because the mass ratio of Fe:O in
wüsite is 1.5 times larger than the
Fe:O ratio in hematite, there must
be 1.5 Fe atoms in a unit of wüsite
and 1 Fe atom in a unit of hematite
Some carbon atoms weigh more
than other carbon atoms
0%
1
0%
0%
2
3
0%
4
The Statements According to Dalton’s Theory are:
4
1. Copper atoms can combine with zinc atoms to make gold atoms
incorrect; atoms of one element cannot turn into atoms of another
element by a chemical reaction
2. Bulk water is composed of many identical molecules, each having
one oxygen atom and two hydrogen atoms
correct; atoms combine together in compounds in small whole
number ratios, so that you could describe a compound by
describing the number of atoms of each element in a molecule.
3. Some carbon atoms weigh more than other carbon atoms
incorrect; all atoms of the same element weigh the same
4. Because the mass ratio of Fe:O in wüsite is 1.5 times larger than
the Fe:O ratio in hematite, there must be 1.5 Fe atoms in a unit of
wüsite and 1 Fe atom in a unit of hematite
incorrect; atoms must combine in small whole-number ratios. You
can get the Fe:Fe mass ratio to be 1.5 if the formula for wüsite is
FeO and the formula for hematite is Fe2O3
Charge
•
•
•
•
Two kinds of charge exist in nature, + and –
Two opposite charges attract, + attracts –
Two like charges repel, + repels +, – repels –
To be neutral, either no charge is present, or
equal amounts of + and – charges are
present
5
Cathode Ray Tube (CRT)
• An evacuated glass tube with metal electrodes
connected to a high voltage power supply
• A glowing ray emanates from the cathode and
terminates at the anode
6
What is the ray of a cathode ray tube?
light or charged particles
• Thomson believed that the cathode ray was
composed of tiny electrically charged particles
He designed an experiment to demonstrate this by
measuring the amount of force it took to deflect the
particles path a given amount
A force was exerted by placing a charged electric
field around the CRT
1. charged matter is attracted to an electric field
2. light’s path is not deflected by an electric field
7
Thomson’s Experiment
+++++++++++
Cathode
Anode
(+)
(-)
-------------
-
Power Supply
8
+
Thomson’s Results
• the beam was deflected toward the + plate, 
 Cathode rays are made of tiny negatively charged
particles, not light
• Every material tested had these same particles
• The amount of deflection is related to two factors,
the charge and mass of the particles
The charge to mass ratio of these particles is
−1.76 x 108 C/g
(high charge/mass ratio)
Compared to the charge to mass of the hydrogen
ion, which is
+9.58 x 104 C/g
(lower charge/mass ratio)
9
Thomson’s Conclusions
• Since atoms are neutral as elements, their +
•
•
and – charges have to cancel
If the – particle has the same amount of
charge as a + hydrogen ion, then the –
particle must have a mass almost 2000 times
smaller than a hydrogen ion!
The only way for this to be true is if these
particles were pieces of atoms
apparently, the atom is not unbreakable
10
Thomson’s Conclusions
• Thomson believed that these particles were
therefore the ultimate building blocks of matter
“We have in the cathode rays matter in a new state,
a state in which the subdivision of matter is carried
very much further . . . a state in which all matter . . .
is of one and the same kind; this matter being the
substance from which all the chemical elements are
built up.”
• The cathode ray particles became known as
electrons
tiny, negatively charged particles found in all atoms
11
Millikan’s Oil Drop Experiment
pinhole
(X-rays)
12
Millikan’s oil droplet experiment
• An oil droplet that falls into the electric field becomes
negatively charged by the stream of ionizing radiation
 When the charge on the drop (in coulombs) / oil drop mass (g)
= acceleration due to gravity / the applied electric field, the
droplet stops falling
q/m = g/E
• The oil drop mass = oil density x oil drop volume
 This was used to calculate the C/g ratio.
 a whole # of electrons was assumed
 every drop had a C/g ratio that reduced to a charge of
-1.60 x 1019 C, the charge of the electron
the electron has a mass of 9.1 x 10-28 g
electrons are particles found in all atoms
13
A New Theory of the Atom
• Because of these experiments, the atom was no
•
longer believed to be indivisible
Thomson proposed a new atomic model that
replaced Dalton’s premise that each element is
composed of tiny, indestructible particles (atoms)
The rest of Dalton’s theory was still valid at this point
• Thomson proposed that instead of being a hard,
marble-like unbreakable sphere, the way Dalton
described it, the atom actually had an inner
structure
16
J.J. Thomson’s Plum Pudding Model
• Atoms contain many negatively charged electrons
held in the atom by their attraction for the atom’s
positively charged electric field
 The mass of the atom is due to the mass of the electrons
 Thomson assumed there were no positively charged
pieces in the atom, because none showed up
in the cathode ray experiment
 the negatively charged particles
should not be near each other
because
they would repel
• Atoms are mostly empty space
17
Discovery of Radioactivity
Marie Curie
1867-1934
• Henri Becquerel and Marie Curie
discovered that certain elements
constantly emit small, energetic particles
and rays that could penetrate matter
• Ernest Rutherford discovered that there were three
different kinds of emissions
 alpha (a) particles with a mass 4x that of the H atom and a
positive (+) charge
 beta (b) particles with a mass ~1/2000x that of the H atom
and a negative (–) charge
 gamma (g) rays that are energy rays, not particles
18
Rutherford’s Experiment to determine
if an atom is mostly empty space
Devised a method to shoot alpha particles through
atoms to see the result
• a particles have a mass of 4 amu & a charge of +2 c.u.
 large compared to electrons, so electrons will not deflect
them
• used large target atoms of gold, mass = 197 amu
 used very thin sheets of gold so the “bullet” would not be
absorbed
19
20
Rutherford’s Experimental Results
and Conclusions
• > 98% of the a particles went straight through
 Atoms are mostly empty space because >98% of the
particles went straight through
• ~ 2% of the a particles went through the foil but
were deflected by large angles
 The dense particle is positively charged which
explains the large deflections of ~2% of the particles
• ~ 0.005% of the a particles bounced off the foil
 Atom contain a dense particle that is small in volume
compared to the atom but large in mass because ~
0.005% of the particles bounced back
21
Plum Pudding
Atom
•
•
•
•
•
•
•
•
•
•
A few of the
a particles
do not go through
•
•
•
•
•
•
•
•
•
•
•
•
If atom was like
a plum pudding,
all the a particles
should go
straight through
Nuclear Atom
.
.
.
Almost all a particles
go straight through
Some a particles
go through, but are deflected due to
+:+ repulsion from the nucleus
22
Rutherford’s Interpretation
The Nuclear Model
1. Atoms contain a tiny dense center, the nucleus
2. The nucleus contains the atom’s entire mass
 the electron’s comparative weight/mass is negligible
3. The nucleus is positively charged and it balances
4.
the negative charges of the electrons
The electrons are dispersed in the empty space
of the atom surrounding the nucleus
23
Drop a pea into any large sports stadium. The pea is to the
nucleus as the stadium is to the atom.
the scale of comparison is called a relative scale
Structure of the Nucleus
• Rutherford proposed that the nucleus had a particle
•
that had the same amount of charge as an
electron, but opposite in sign, called protons
Protons are found in the nucleus and have a
charge=+1.60 x1019 C and mass=1.67262 x 10−24 g
 Each proton has a charge of +1 charge units (cu)
 protons and electrons have equal amounts of charge, but
opposite in sign, therefore
 atoms must have equal numbers of protons and
electrons to be neutral
25
Something is Missing
• How can an atom’s nucleus have multiple protons
with + charges placed in close proximity?
 they should repel each other
• Beryllium atoms have four protons
 it should weigh 4 amu; but it actually weighs 9.01 amu
the protons account for only 4 amu
Be has four electrons which weigh 4x(~0.00055 amu), far
less than the extra 5 amu needed for mass balance
 Where is the extra mass coming from?
26
Proposal to Explain the Difference
• Rutherford and Chadwick proposed that there
•
•
•
was another particle in the nucleus – called a
Neutron
Neutrons are subatomic particles with a mass
= 1.67493 x 10−24 g
1 amu (slightly heavier than a proton)
They have no charge
They are found in the nucleus
27
Atomic Mass Units
• We take 1/12th the mass of the carbon atom (6 protons
and 6 neutrons), and call it 1 atomic mass unit (amu)
 protons and neutrons have a mass of just over 1 amu
 electrons have a mass of 0.00055 amu
too small to be relevant to the total mass of the atom
28
29
Why Does Matter Appear Continuous
If the Atom Is Mostly Empty Space?
The emptiness of the atom is on such a
small scale that the variations in density
cannot be seen
Consider a ladder framework that is
mostly empty (right). This framework, if
large enough, will look solid from a
distance. Also, it has some rigidity based
on the interaction of adjacent latices.
30
Elements
• Each element has a unique number of protons
in its nucleus
The number of protons in the nucleus is called the
atomic number, “Z”
The number of protons defines the element
• All the elements are arranged on the Periodic
Table in order of their atomic numbers
Each element has a unique name and symbol
Each symbol is either one or two letters
– one capital letter alone, or
– one capital letter followed by one lowercase letter
31
Some symbols come from the element ‘s name, like C for carbon. Others come from
the Latin name, like gold (Au-aurum) and copper (Cu-cuprium), or the Greek, like
chlorine (Cl-chloros) and argon (Ar-argos).
The Periodic Table of the Elements
The atomic number tells you
how many protons are in the
nucleus and how many
electrons are in the atom
32
The Nucleus
• Atomic number (Z) = # of protons
 “Z” does not change for each element
 Does change from element to element
• Mass Number (A) = # protons + # neutrons
 “A” does change for elements and their isotopes
• Both Z and A are whole numbers
33
Isotopes
• It was discovered that the same element could
have atoms with different masses called
isotopes
• All isotopes of an element have the same number
•
of protons and electrons
All isotopes of an element are chemically identical
undergo the exact same chemical reactions
• Isotopes of an element have different masses
 Due to different numbers of neutrons
 Isotopes are identified by their mass numbers, the
sum of all the protons and neutrons in the nucleus
34
Isotopes
• The observed atomic mass of any element is
a weighted average of the weights of all the
naturally occurring isotopes
 there are two isotopes of chlorine found in nature, one
that has a mass of about 35 amu and another that
weighs about 37 amu
 the percentage of each isotope is called the isotope’s
natural abundance
 the observed atomic mass of chlorine is 35.45 amu
35
Neon
Symbol
Number of Number of A, Mass
Protons
Neutrons Number
Percent
Natural
Abundance
Ne-20 or 20
10Ne
10
10
20
90.48%
21Ne
Ne-21 or 10
10
11
21
0.27%
22Ne
Ne-22 or 10
10
12
22
9.25%
36
Which is not an example of isotopes?
1.
2.
3.
4.
Z=1, N=1 and Z=1, N=2
Z=9, N=9 and Z=8, N=9
Z=9, N=9 and Z=9, N=10
Z=1, N=3 and Z=1, N=2
0%
37
1
0%
2
0%
3
0%
4
Example 2.3b: How many protons, electrons,
52 Cr
and neutrons are in an atom of 24
?
Given:
Find:
Conceptual
Plan:
52
24 Cr
therefore A = 52, Z = 24
# protons, # electrons, # neutrons
atomic
number
symbol
symbol
atomic number
& mass number
# p+
# e−
# n0
in neutral atom, # p+ = # eRelationships:
mass number = # protons + # neutrons
0
A
=
Z
+
#
n
+
−
Solution: Z = 24 = # p = # e
52 = 24 + # n0
52-24 = # n0
28 = # n0
Check:
for most stable isotopes, n0 ≥ p+
38
Complete the table and add all the numbers in the
Electrons column. They add up to ?????
27
13 Al
1.
2.
3.
4.
122
116
106
98
0%
39
1
0%
0%
0%
2
3
4
Complete the table
13
6C
96
42 Mo
27
13 Al
133
55 Cs
116
40
Reacting Atoms
• Elements that undergo chemical reactions do not turn
into other elements
 All the atoms present at the start will be there at the end
 The number of protons determines the element
 the number of protons don’t change in a chemical reaction
• Reactions involve sharing or transfer of electrons
• When atoms gain or lose electrons, they acquire a
charge to become ions
 Atoms that gain electrons become negatively charged
anions
 Atoms that lose electrons become positively charged
cations
41
Ions and Compounds
• Ions behave differently than neutral atoms
metallic sodium, made of neutral Na atoms, is
highly reactive and quite unstable
sodium cations, Na+, are found in table salt and
are very nonreactive and stable
• Because materials such as table salt are
neutral overall, there must be a balance
between the charges of cations and anions
42
Atomic Structures of Ions
• Nonmetals form anions
• For each negative charge, the ion has
one more electron than the neutral atom
F = 9 p+ and 9 e−
F− = 9 p+ and 10 e−
P = 15 p+ and 15 e−
P3− = 15 p+ and 18 e−
Anions are named by changing the ending of
the element’s name to -ide
fluorine
oxygen
F + 1e−  F−
O + 2e−  O2−
43
fluoride ion
oxide ion
Atomic Structures of Ions
• Metals form cations
• For each positive charge, the ion has one less
electron than the neutral atom
Na atom = 11 p+ and 11 e− Na+ ion = 11 p+ and 10 e−
Ca atom = 20 p+ and 20 e− Ca2+ ion = 20 p+ and 18 e−
• Cations are named the same as the
elemental metal
sodium
calcium
Na  Na+ + 1e− sodium ion
Ca  Ca2+ + 2e− calcium ion
44
Complete the table, add the ion charge column, and provide
your answer.
A.
B.
C.
D.
E.
+2
-2
0
+3
-3
Al 3 
0%
45
A.
0%
0%
B.
C.
0%
0%
D.
E.
Complete the table
S
2
Mg 2 
Al 3 
Br 
46
Mendeleev
• Periodic Law – Saw a repeating pattern of
properties when the elements are arranged in
order of increasing atomic mass
• Ordered elements in rows by increasing atomic mass
from left to right
• Put elements with similar properties in the same
column, with heavier mass lower in the column
• Used pattern to predict properties of undiscovered
elements
47
Periodic Pattern
48
Most
About
A
fewofelements
¾the
of remaining
the elements
are classified
elements
are classified
asare
metalloids.
classified
as metals.
as
nonmetals.
They have
Their
solidsaTheir
have
reflective
solids
somesurface,
characteristics
have a non-reflective
conductofheat
metals
and
surface,
electricity
and
some
dobetter
of
not
nonmetals.
conduct
than other
heatelements,
and electricity
and are
well,
and
malleable
are brittle.
and ductile
49
Metals
• All metals are solids at room temperature
 except Hg which is liquid
•
•
•
•
•
•
•
Have reflective, shiny surface
Conduct heat & electricity
Are malleable (can be shaped)
Are ductile (can be drawn or pulled into wires)
Lose electrons to form cations in reactions
Comprise ~75% of the elements
Found on the lower left of the periodic table
50
Nonmetals
Sulfur, S(s)
• Found in all three states
• Poor conductors of heat &
•
•
•
electricity
Are brittle
Gain electrons in reactions to
become anions
Found in upper right quadrant
of the Periodic table
Hydrogen, a nonmetal) is an
exception
51
Bromine, Br2(l)
Chlorine, Cl2(g)
Metalloids
• Show some properties of metals and some of
•
nonmetals
Also known as semiconductors
Properties of Silicon
shiny
conducts electricity
does not conduct heat well
brittle
52
The Modern Periodic Table
• Columns are called Groups
Elements with similar chemical and physical
properties are in the same group
designated by a number and letter at top
• Rows are called Periods
Each period shows a pattern of properties
repeated in the period above or below
53
The Modern Periodic Table
• Main group elements have letter “A” designation
• Transition elements have letter “B” designation
All transition elements are metals
• The 2 bottom rows are called inner transition
elements
aka rare earth elements
are metals
really belong in Period 6 & 7
54
55
= Alkali metals
= Halogens
= Alkali earth metals
= Lanthanides
= Noble gases
= Actinides
= Transition metals
56
Hydrogen
• Nonmetal
• Colorless, diatomic gas
 very low melting point, boilng point and density
• Reacts with nonmetals to form molecular
compounds
 Reacts with Cl2 to form HCl, an acidic gas
 HCl dissolves in water to form acids
 Reacts with O2 to form H2O, a liquid
• Reacts with metals to form hydrides
 metal hydrides react with water to form H2
• Placed in Group IA
 but based on properties, it does not belong there
57
Group IA = Alkali Metals
•
•
•
•
Low density
Low melting point
Soft
Very reactive
lithium
 never found in elemental form
• React with water to form basic
sodium
(alkaline) solutions and H2
2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g)
 releases a lot of heat
• Tend to form water-soluble salts
• Salts isolated from seawater are
•
melted and electrolyzed to form
elemental form
Flame tests  Li = red, Na =
yellow, K = violet
58
potassium
rubidium
cesium
Group IIA = Alkali Earth Metals
• Harder, higher melting, and denser metals than
•
•
alkali metals
Reactive, but less reactive than alkali metals
Form stable, insoluble oxides
 Metals are extracted from the oxides
• Oxides are basic (alkaline earth)
• Like IA, IIA metals react with water to form H2
 Be = no rxn (NR); Mg = rxn with steam; Ca, Sr, Ba = rxn
with cold water
“rxn” abbreviation for reaction
• Flame tests  Ca = red, Sr = red, Ba = green
59
Group VIIA = Halogens
• Nonmetals
 Very reactive
 React with metals to form ionic
compounds
 Cl2, Br2 react slowly with water
Br2 + H2O  HBr + HOBr
• All diatomic (X2)
fluorine
chlorine
bromine
 F2 and Cl2 are gases
 Br2 liquid
 I2 solid
iodine
• HX all acids
astatine
HF weak < HCl strong< HBr < HI
60
Group VIIIA = Noble Gases
• All gases at room temperature
 very low melting and boiling points
• Very unreactive, practically inert
• Very hard to remove an electron from or
give an electron to a noble gas
61
Ion Charge and the Periodic Table
• The charge on an ion can often be determined
•
from an element’s position in the Periodic Table
Metals always form positively charged species,
cations
For many main group metals, the charge equals the
group number
• Nonmetals form negatively charged species,
anions
the charge equals the group number
62
63
potassium cation
sulfide anion
calcium cation
bromide anion
aluminum cation
Group IA
Group VIA
Group IIA
Group VIIA
Group IIIA
K+
S2−
Ca2+
Br−
Al3+
Counting Atoms
• a mole is the mass of a material that contains
6.022 x 1023 atoms/particles/ things
Avogadro’s Number
• the mole’s value is referenced to the number of
atoms in exactly 12 grams of pure carbon-12
 1 atom of C-12 weighs exactly 12 amu
 the average atomic mass of C is12.01 amu, therefore
 1 mole of C atoms weighs 12.01 g
 12.01 g contains 6.022 x 1023 atoms
• this provides a relationship between mass and the
number of atoms
64
Counting Atoms by Moles
The number of atoms, 6.022 x 1023,
per mole can be used as a conversion
factor. Given the number of moles of
a substance, you can calculate the
number of atoms, and visa versa
65
Example 2.6: Calculate the number of
atoms in 2.45 mol of copper
Given:
Find:
Conceptual
Plan:
Relationships:
2.45 mol Cu
atoms Cu
mol Cu
atoms Cu
1 mol = 6.022 x 1023 atoms
Solution:
Check: because atoms are small, the large number of
atoms makes sense
66
A silver ring contains 1.10 x 1022 silver atoms. How many
moles of silver are in the ring?
1 mol = 6.022 x 1023 atoms
1. 6.62x1045 moles
2. 0.0183 moles
3. 54.7 moles
0%
67
1
0%
2
0%
3
A silver ring contains 1.1 x 1022 silver atoms. How
many moles of silver are in the ring?
Given: 1.1 x 1022 atoms Ag
Find: moles Ag
Conceptual
Plan:
Relationships:
atoms Ag
mol Ag
1 mol = 6.022 x 1023 atoms
Solution:
Check: because the number of atoms given is less than
Avogadro’s number, the answer makes sense
68
Relationship Between Moles and Mass
• Molar Mass = the mass of 1 mole of atoms
6.022x1023 particles
• The molar mass of an element, in grams, is
numerically equal to the element’s atomic mass,
in amu
the lighter the atom, the less a mole weighs
 the more atoms there are in 1 g
• Molar mass can be used as a conversion factor
to convert mass to moles: g x mol/g = mole
• or moles to mass: mol x g/mol = g
69
Mole and Mass Relationships
1 mole
carbon
12.01 g
1 mole
sulfur
32.06 g
70
Example 2.7: Calculate the moles of carbon
in 0.0265 g of pencil lead
Given: 0.0265 g C
Find: mol C
Conceptual
Plan:
gC
mol C
Relationships: 1 mol C = 12.01 g
Solution:
Check:
because the given amount is much less
than 1 mol C, the number makes sense
71
Calculate the moles of sulfur in 57.8 g of sulfur
1 mol S = 32.07 g
A.
B.
C.
D.
1.80
0.555
1850
1.802
0%
A.
72
0%
0%
B.
C.
0%
D.
Calculate the moles of sulfur in 57.8 g of sulfur
Given:
Find:
Conceptual
Plan:
57.8 g S
mol S
gS
mol S
Relationships: 1 mol S = 32.07 g
Solution:
Check: because the given amount is much less than 1
mol S, the number makes sense
73
Example 2.8: How many copper atoms are in
a penny weighing 3.10 g?
Given:
Find:
Conceptual
Plan:
Relationships:
3.10 g Cu
atoms Cu
g Cu
mol Cu
atoms Cu
1 mol Cu = 63.55 g, 1 mol = 6.022 x 1023
Solution:
Check:
because the given amount is much less
than 1 mol Cu, the number makes sense
74
How many aluminum atoms are in a
can weighing 16.2 g?
Given: 16.2 g Al
Find: atoms Al
Conceptual
Plan:
g Al
mol Al
atoms Al
Relationships: 1 mol Al = 26.98 g, 1 mol = 6.022 x 1023
Solution:
Check: because the given amount is much less than 1
mol Al, the number makes sense
75
Mass Spectrometry
• Mass and isotope abundance are measured with
•
a mass spectrometer
Atoms or molecules are ionized, then accelerated
down a tube
some molecules are broken into fragments during
the ionization process
these fragments can be used to help determine the
structure of the molecule
• Their path is bent by a magnetic field, separating
them by mass
76
Mass Spectrometer
77
Mass Spectrum
• A mass spectrum is a
chart that gives the
relative mass and
relative abundance of
each particle present
78
Example 2.5: If copper is 69.17% Cu-63 with a mass of
62.9396 amu and the rest Cu-65 with a mass of 64.9278 amu,
find copper’s atomic mass
Given:
Find:
Conceptual
Plan:
Relationships:
Cu-63 = 69.17%, 62.9396 amu
Cu-65 = 100-69.17%, 64.9278 amu
atomic mass, amu
isotope masses,
isotope fractions
avg. atomic mass
Solution:
Check:
the average is between the two masses,
closer to the major isotope
79
Practice – Ga-69 with mass 68.9256 amu and abundance of
60.11% and Ga-71 with mass 70.9247 amu and abundance of
39.89%. Calculate the atomic mass of gallium.
Given: Ga-69 = 60.11%, 68.9256 amu
Ga-71 = 39.89%, 70.9247 amu
Find: atomic mass, amu
Conceptual
isotope masses,
avg. atomic mass
Plan:
isotope fractions
Relationships:
Solution:
Check:
the average is between the two masses,
closer to the major isotope
80
81
Scanning Tunneling Microscope
• Gerd Bennig and Heinrich
Rohrer found that as you
pass a sharp metal tip over a
flat metal surface, the
amount of current that flows
varies with distance between
the tip and the surface
• Measuring this “tunneling”
current allowed them to scan
the surface on an atomic
scale – essentially taking
pictures of atoms on the
surface
82
Operation of a STM
83
Scanning Tunneling Microscope
• Later scientists
found that not only
can you see the
atoms on the
surface, but the
instrument allows
you to move
individual atoms
across the surface
84