Transcript Chemistry 2: matter is made up of atoms

Chemistry 2: Scientific
Measurement
• Quantitative measurements give results
in a definite form, usually as numbers
• Qualitative measurements give results in
a descriptive nonnumeric form
Accuracy and Precision
• Accuracy is how close a measurement
comes to the actual dimension or true
value of whatever is measured
• Precision is concerned with the
reproducibility of the measurement
Significant Figures in
Measurements
• The significant figures in a measurement
include all the digits than can be known
precisely plus a last digit that must be
estimated.
• 1. Every nonzero digit is significant.
• 2. Zeros appearing between nonzero digits are
significant.
• 3. Zeros appearing in front of all nonzero digits
are not significant.
• 4. Zeros at the end of a number and to the
right of a decimal point are significant.
• 5. Use scientific notation whenever possible.
The Metric System
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Length
Mass
Time
Electric current
Temperature
Amt of substance
Light intensity
Pressure
Energy
Force
Capacitance
meter
kilogram
second
ampere
kelvin
mole
candela
pascal
joule
newton
farad
m
kg
s
A
K
mol
cd
pa
J
N
F
Units of Length, meter
• kilo
• deci
• centi
• milli
• micro
• Nano
• Pico
k
d
c
m
m
n
p
1000
1/10
1/100
1/1000
103
10-1
10-2
10-3
10-6
10-9
10–12
Units of Volume, liter
• A liter is the volume of a cube that is 10
cm on each edge
(10cmx10cmx10cm=1000cm3=1 L
Units of Mass, kilogram
• A kilogram is the mass of 1 L of
water at 4 oC.
• A gram is defined as the mass of 1
cm3 of water at 4 oC.
Density and Specific Gravity
• Density is the ratio of the mass of an
object to its volume.
• Density = mass / volume
• Specific gravity is a comparison of the
density of a substance to the density of a
reference substance, usually at the same
temperature.
• Specific gravity = density of substance
(g/cm3) /density of water (g/cm3)
Measuring Temperature
• Temperature is the degree of hotness or
coldness of an object.
• Heat transfer occurs when two objects at
different temperatures contact each other.
• Swedish astronomer Anders Celsius (17011744)—the Celsius scale takes the freezing
point of water as 0 oC and boiling point at 100
oC at one atmosphere pressure.
• Scottish physicist Lord Kelvin (1824-1907)—the
Kelvin scale, the freezing point of water is 273
K, and the boiling point is 373 K.
• The zero point (0 K) on the Kelvin scale is
absolute zero. It is –273 oC.
Measuring Heat
• English physicist James Joule (18181889) formulated heat conversion.
• 1 J = 0.239 cal and 1 cal = 4.18 J
• 1 calorie is the quantity of heat that
raises the temperature of 1 g of pure
water 1 oC.
• 1 kcal = 1000 cal = 1 Cal.
Specific Heat Capacity
• The quantity of heat required to
change an object’s temperature by
exactly 1 oC is the heat capacity of
that object.
• The specific heat capacity, or simply
the specific heat, of a substance is
the quantity of heat required to raise
the temperature of 1 g of the
substance 1 oC.
Matter is made up of atoms
• Joseph Proust, in 1799, observed
that water consists of 11% hydrogen
and 89% oxygen: in a ratio of 1:8 by
mass
• Antoine Lavoisier, in 1774,
discovered the law of conservation of
matter
• John Dalton in 1800s proposed the
atomic theory
Dalton’s Atomic Theory
• 1. All matter is made up of atoms
• 2. Atoms are indestructible and
cannot be divided into smaller
particles (not true anymore)
• 3. All atoms of one element are
exactly alike, but they are different
from atoms of other elements
The Methods of Science
• Repeated observation give rise to
hypothesis, which is tested by
experiments
• Repeated experiments either confirm the
hypothesis or revise the hypothesis
• After a hypothesis has been verified my
other scientists, it becomes a theory
• More experiments will give rise to a
revised theory
• When the theory is firmly established
without exceptions, it becomes a law
The Discovery of Atomic
Structure
• J.J. Thomson, in 1897, discovered the
electron using the cathode-ray tube, and
that electrons are negatively charged
• Rutherford’s gold foil experiments showed
that atoms have positively charged nuclei
• Thomson’s atomic model: electrons
embedded in a ball of positive charge
• Nagaoka’s atomic model: electrons orbit
around the positive nucleus like planets
round the sun
Atomic Numbers and
Masses
• Atomic number is the number of protons in
the nucleus of an element
• Mass number is the sum of protons and
neutrons in the nucleus. It is the average of
all isotopes of an element
• Isotopes are different forms of an element
having the same number of protons but
different number of neutrons, hence their
masses are different
• Atomic mass, 1 u = 1/12 the mass of C-12
Common Particles of an
Atom
• Particle, symbol, charge, Z, mass in u
• Proton
p+ 1+
1
1.01
• Neutron
n0
0
1
1.01
• Electron
e10
0.00055
Electrons in Motion
• Niels Bohr proposed that electrons have
energies to move round the nucleus
• Now electrons are considered as particles
and waves
• Electrons move round the nucleus near
the speed of light
• Electrons cannot be accurately located,
only the probability of finding it
The electromagnetic
spectrum
• All forms of radiant energy can be
placed in the electromagnetic
spectrum, from radio waves, AM, FM,
microwaves, infrared, visible
spectrum, UV, X-rays, to gamma
rays, which are very energetic and
have very short wave-lengths
Electron and Light
• Hydrogen can be shown to consist of
many energy levels, which are
similar to the rungs on a ladder
• The electron cloud model of an atom
suggests that energy levels are
concentric spherical regions of space
around the nucleus
• Electrons in the outermost energy
level are called the valence electrons
Lewis Dot Diagrams
• A Lewis dot diagram illustrates
valence electrons as dots around the
chemical symbol of an element. Each
dot represent one valence electron
• Chemical changes only involve the
outermost electrons of an element