Science Focus 10
Unit 1 Energy and Matter in Chemical
Change p. 2-137
In this unit we will explore:
• Atoms, elements and compounds
• Ionic, molecular, acidic and basic compounds
• Naming chemical compounds
• The mole
• Water
• Chemical equations and reactions
•What-Is-the-Importance-of-Chemistry-
Chemical Composition of the Human Body
•
This is the elemental chemical composition of the average
adult human body. Water is the most abundant chemical
compound in living human cells, accounting for 65-90% of each
cell.
•
Each water molecule consists of two hydrogen atoms
bonded to one oxygen atom, but the mass of each oxygen atom
is much higher than the combined mass of the hydrogen.
•
All organic compounds contain carbon, which is why
carbon is the second most abundant element in the body. Six
elements account for 99% of the mass of the human body:
oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus.
Although aluminum and silicon are abundant in the earth's
crust, they are found in trace amounts in the human body.
•
•
•
•
•
•
•
•
•
•
•
•
•
Elements in the Human Body
Element Percent by Mass
Oxygen 65
Carbon 18
Hydrogen 10
Nitrogen 3
Calcium 1.5
Phosphorus 1.2
Potassium 0.2
Sulfur 0.2
Chlorine 0.2
Sodium 0.1
Magnesium 0.05
• Iron, Cobalt, Copper, Zinc, Iodine <0.05 each
• Selenium, Fluorine
<0.01 each
• Value of all this great stuff : less than $1.00
Chapter 1: Atoms, Elements and
Compounds p. 2-39
• 1.1 Working with Chemicals p. 6-10
• Aboriginal peoples have been using chemical
substances for thousands of years to preserve food,
treat illness, build tools and decorate clothing. Many
of these traditional processes are still used today.
tea rich in Vitamin C, also
used to treat stomach and
kidney complaints
treat chest colds and
heart disease
boiled to treat
diarrhea
ASA- asprin; fever
and heart
treatments
7
MSDS
8
MSDS
• Many chemicals, even household products,
have dangerous properties and must be
handled properly. An MSDS, or Material
Safety Data Sheet, lists important information
including physical properties (ie melting and
boiling points), chemical dangers, and how to
store and dispose of the chemical. See figure
1.3 p7
9
10
• WHMIS Symbols:
• The Workplace Hazardous Materials Information System is a
system of warning symbols designed to protect people who
use harmful substances at work.
Oxidizing material
12
• Compressed gas
13
• Class D1: Poisonous and Infectious
Material Causing Immediate and
Serious Toxic Effect
14
• Poisonous and Infectious Material
Causing other Toxic effects
15
• Dangerously Reactive material
16
• Biohazardous Infectious Material
17
• Corrosive Material
18
•Read “Safety in the Chemistry Laboratory”
p. 7-8
•Do BLM 1-1 Interpreting an MSDS
19
• Classifying Matter
• Matter is anything that has mass and occupies space.
– Mixtures can be mechanical (heterogeneous) where the separate
parts are visible (concrete) or solutions (homogeneous) where the
different parts are not visible. (air, apple juice)
– Pure substances include elements (such as Na or Cl) or compounds
(such as NaCl). Compounds (sugar) can be chemically separated into
simpler substances, elements cannot.
Matter
• Pure Substance
Element
Compound
Mixture
Heterogeneous
(Mechanical)
Homogeneous
(Solution)
Pure Substances
• Element- cannot be chemically broken down
into simpler substances
• Compound- two or more elements that are
chemically combined; can be separated
chemically into simpler substances
22
Mixture
• Heterogeneous- different components of the
mixture are visible; composition is variable
throughout the mixture
• Homogeneous- different components are not
visible; composition is constant throughout the
mixture
23
• See figure 1.5 p10.
• Do Practice Problems #1-4 p10
• Do Check Your Understanding #1-3 p11
24
• 1.2 Developing Atomic Theories p.12-24
• The idea that matter is made up of small particles is
rooted in work of scientists many centuries ago.
• One of the earliest theories on matter is credited to
Democritus (~400 BC), who stated that matter is
made up of infinite tiny, indivisible, constantly
moving units. This theory has evolved over recent
history….
Dalton’s “Billiard Ball” Model
• 1766-1844
• Proposed to explain the interaction of chemicals
• Key features:
• Matter is composed of small indivisible particles (atoms) that can
be neither created nor destroyed.
• All atoms of the same element are identical in mass and size, but
different from atoms of other elements.
• Atoms exist in an otherwise empty space and are in constant
motion.
• They may collide to form new combinations (compounds).
• Chemical reactions change the way atoms are arranged, but do not
change the atoms themselves.
• Why model was rejected or modified:
• contained inaccuracies regarding the relative masses of several
atoms
• could not account for the electric nature of matter
Thomson’s “Raisin Bun” model
• 1856-1940
• proposed to explain negative charges in the
atom
• Main Features:
• The atom is a positively charged sphere in which
negatively charged electrons are embedded, like
raisins in a bun.
• Overall, the atom has no charge.
•
Electrons are extremely small
compared with the size of an atom.
28
• Thompson’s Experiment
• Thomson conducted a series of experiments with cathode ray
tubes leading him to the discovery of electrons. (see figure
1.10 p. 15)
• Thomson discovered that cathode rays (a stream of negatively
charged particles) could be bent towards a positively charged
plate.
• He concluded that all atoms contain electrons (small negative
charges) that are distributed throughout a positively charged
solid.
+
• Thompson model was rejected because it
could not explain how radioactive materials
emit alpha particles
30
• Rutherford’s “Solar System” Model
• 1871-1937
• proposed to account for deflection of alpha particles
(+) from gold foil
• Main Features:
• At the centre of every atom is a small, positively charged nucleus.
• The nucleus accounts for the majority of the mass of the atom.
• Electrons are attracted to the nucleus and orbit in a cloud around
the nucleus.
• A third subatomic particle, the neutron, with no charge but a
mass similar to a proton also exists in the nucleus.
• Rejected because electrons do not orbit randomly
• Rutherford’s Experiment:
• Rutherford shot positively charged alpha particles through thin
gold foil. Instead of traveling straight through, some of the
particles were deflected.
• He suggested that the atom must be composed of a very small
core of positively charged particles called the nucleus that is
1/10 000th of the size of the atom.
• It was the nucleus that deflected the positive particles.
Read P. 16-17
The Bohr Model
• 1885-1962
• Proposed to organize electrons
• electrons do not orbit the nucleus randomly, but rather in
specific circular orbits, called energy levels or electron shells
• Transitions of electrons to higher energy levels require
energy.
• Transitions to lower energy levels produce electromagnetic
radiation (light and radio waves).
• Each energy level has a fixed maximum number of electrons
that can reside in it.
• Rejected because at higher levels there are sublevels
33
• The first shell can hold only 2 e• The second and third shell can hold 8 e• The outermost shell is called the valence shell
• Quantum mechanics research has found that
electrons exist in a charged cloud around the
nucleus
Awesomeness….
MIND BLOW!!!!!!!!!!
• http://www.youtube.com/watch?v=DZGINaRU
EkU
36
In Summary
• Dalton’s Billiard Ball : Matter is made up of solid
spheres called atoms
• Thomson’s Raisin Bun: Electrons are embedded in the
atom
• Rutherford: The nucleus is composed of protons and
neutrons. Electrons are outside of the nucleus.
• Bohr: Electrons exist in orbitals/ energy levels
• Quantum Mechanics Theory: Most current accepted
atomic theory.
• atomic theories
38
Features of a Simplified Modern
Model of the Atom
• a tiny, dense nucleus that is surrounded by electrons (e-)
• nucleus contains protons (p+) and neutrons (no), called
nucleons (exception: H-1 nucleus contains one proton
only)
• nucleus accounts for most of the mass of the atom
• e- exist at certain allowed energy levels
• p+ carry a positive charge, e- carry a negative charge and
no carry no charge. See table 1.2 p 22.
• A neutral atom always has equal numbers of e- and p+
• If a nucleus were the size of a baseball, the
electron orbitals around the nucleus would
take up the space of an arena!
Did You Know?
• QUARKS
• Scientists now believe that neutrons and
protons are made up of even smaller particles
called quarks. They believe that matter is
made up of dozens more of these subatomic
particles.
• . . . what do you remember about the periodic
table?
The Periodic Table tells us Many Things About
the Elements
• 11
23
•
Na
• sodium
•
atomic mass
atomic number
symbol
name of element
• Atomic Number
• # of protons in the nucleus
• Also indicates the # of electrons (since all atoms
are neutral in charge).
• (Note: the atomic number increases by 1 from
left to right across the periodic table.)
• Atomic Mass
• total number of protons and neutrons in the
nucleus.
• Atomic mass is an average mass – it varies
according to how many neutrons are present in
each atom (isotopes).
• Mass is measured by atomic mass units (amu).
• Protons and neutrons each have an amu of 1
• How many protons are there in an
oxygen atom?
• How many electrons?
• How many neutrons?
• How many protons are there in an
oxygen atom? 8
• How many electrons?
8
• How many neutrons?
16 – 8 = 8
• How many protons are there in a silicon
atom?
• How many electrons?
• How many neutrons?
• How many protons are there in a silicon
atom? 14
• How many electrons? 14
• How many neutrons? 28 – 14 = 14
• How many protons are there in a nitrogen
atom?
• How many electrons?
• How many neutrons?
• How many protons are there in a nitrogen
atom? 7
• How many electrons? 7
• How many neutrons? 14 – 7 = 7
• Isotopes and Nuclear Notation
• Isotopes of an atom have the same number of protons, but different
numbers of neutrons and therefore different atomic masses.
• The masses reported on the periodic table are weighted averages of
all naturally occurring isotopes.
• Scientists use the following notation to describe a specific isotope:
•
• mass = total number of protons + neutrons
• atomic number = total number of protons in the nucleus
• These two numbers can be used to determine the number of
neutrons:
• mass number – atomic number = # neutrons
• eg 1 . Silicon-29
or
• For Silicon-29, the number of neutrons is 29-14=15!
• eg 2. the most abundant isotope of uranium is uranium-238. The
complete symbol of this isotope is:
• How many protons in uranium-238 ?
• …electrons?
• …neutrons?
• How many protons are in uranium-238 ?
• All Uranium atoms have 92 protons! That’s
why Uranium is #92 on the periodic table!
• How many protons are in uranium-238 ?
• All Uranium atoms have 92 protons! That’s
why Uranium is #92 on the periodic table!
• How many electrons?
• How many protons are in uranium-238 ?
• All Uranium atoms have 92 protons! That’s
why Uranium is #92 on the periodic table!
• How many electrons?
• 92 of course!
• How many protons are in uranium-238 ?
• All Uranium atoms have 92 protons! That’s
why Uranium is #92 on the periodic table!
• How many electrons?
• 92 of course!
• How many neutrons?
• How many protons are in uranium-238 ?
• All Uranium atoms have 92 protons! That’s
why Uranium is #92 on the periodic table!
• How many electrons?
• 92 of course!
• How many neutrons?
• mass number – atomic number = # neutrons
• so with uranium-238, the number of neutrons
is 238-92=146!
Isotopes of Hydrogen
hydrogen-1
hydrogen-2
hydrogen-3
Radioactive Isotopes
-Isotopes that have unstable nuclei and “fall apart” giving off
radiation
Cobalt – 60 is used to kill cancer cells in radiation
therapy
Uranium-235 is used to make atomic bombs!
• Read Nuclear Notation p.22
• Do Practice Problems p. 23 #5-8
• Do Check Your Understanding p. 24 #1-4
1.3 Electrons and the Formation of Compounds p. 25-37
• During the 1800’s, a Russian chemist, Dmetri Mendeleev,
examined 62 elements. He developed a table of these elements
based upon their repeating properties.
• He also predicted the existence and
properties of unknown elements and left
spaces on the periodic table for them.
• Elements are arranged according to
increasing atomic number (number of
protons in the nucleus)
65
• Examine our modern periodic table on p482. It
displays the known elements in a format that follows
various patterns and trends:
• Patterns in the Periodic Table
• Left of the staircase line are metals: mostly solids,
these elements are shiny, malleable, ductile, and
conduct electricity.
• To the right are non-metals: these elements are
either solid, liquid or gas. They are dull, brittle and
do not conduct electricity.
• Surrounding this line are the metalloids: these
elements display both metal and non-metal
properties.
• Periods (rows)
• The period number tells you how many energy levels
you have.
• Properties change in 2 ways as you move from left to
right across a period:
– The elements change from metal to non-metal
– The elements become less reactive.
• Groups
• Elements in the same group have very similar
properties.
• Group number tells us how many electrons are in the valence
shell.
• for groups 13 – 18, we use the last number to designate the
number of valence electrons
• (eg. elements in group 16 have 6 valence electrons)
• electrons fill the first orbital before they can occupy the
second, and fill the second before they can occupy the third
• when the valence level is full, it is referred to as a stable octet
since there are 8 electrons occupying the orbital (unless it is
the first level)
• Group 1 – hydrogen and the alkali metals.
• The most reactive metals and react violently in
air or water. Reactivity increases as you move
down the group
alkali metals and water
Brainiacs show alkali reactions
francium in water
• Group 2 –the alkaline-earth metals.
• very reactive with oxygen, but less reactive
than the alkali metals.
• Group 17 –the halogens
• The most reactive of the non-metals. They
tend to combine with other elements to make
compounds.
• Group 18 – the noble gases
• The most stable and unreactive of the
elements.
•
Inner-transitional elements
• top period = lanthanoids (fits in period 6)
• bottom period = actinoids (fits in period 7)
• Now here’s a little song about the
elements.
• "The Elements". A Flash
animation
Enjoy!
Energy Level Diagrams
• Recall that Bohr inferred that electrons orbit the nucleus in
fixed energy levels.
• Each level can only hold a certain maximum number of
electrons.
• The first can hold 2 electrons, the second can hold 8, and the
third can hold 8.
• Electron energy level diagrams show us the number of electrons in
each energy level, the number of protons, and the charge on the
atom or ion.
•
•
Ex. The energy level diagram for Mg
•
•
2e•
8e•
2e•
12p+
•
•
Mg
•
• period number shows the number of orbitals used in each
element (eg. period 2 elements have 2 orbitals)
• group number describes how many electrons are found in the
valence or outermost energy level (eg. lithium is in group 1 and
has 1 valence electron)
• for groups 13 – 18, we use the last number to designate the
number of valence electrons (eg. elements in group 16 have 6
valence electrons)
• electrons fill the first orbital before they can occupy the
second, and fill the second before they can occupy the third
• when the valence level is full, it is referred to as a stable octet
since there are 8 electrons occupying the orbital (unless it is
the first level)
• The diagram representing the element
beryllium looks like this:
•
2e•
2e•
4p+
•
4 no
• - Draw the energy level diagram for fluorine
atom and a fluoride ion.
• See figure 1.22B on P. 26
• Do Practice Problems p. 27 #9-12
•
and BLM 1-4 “Periodic Table Scavenger
Hunt”.
Electron Dot Diagrams, a.k.a. Lewis Dot Diagrams
• American chemist G.N. Lewis invented these structures to
visualize and track electrons during bond formation.
• To draw:
1) write its chemical symbol
2) surround it by dots that represent the atom’s
valence electrons (use the group number to determine the
number of valence electrons)
3) if an atom has more than 4 valence electrons, the
additional electrons are paired.
4) elements in the same group will have identical dot
diagrams since they have the same number of valence e-
• Examples:
89
Lewis Dot Diagrams of Ions
• Lab: Classifying Matter
• Do BLM 1-5 and then BLM 1-6 on Electron Dot
Diagrams.
Formation of Ions
• Any atom or group of atoms that either loses or gains eand carries either a positive or negative charge is called an
ion.
• Cation - positive charge (has fewer e- than p+)
- metallic atoms lose electrons, eg. Na+
- remember: cats have paws (pos)
• Anion - negative charge (has more e- than p+)
- non-metallic atoms tend to gain electrons
from other atoms, eg. Cl• See p29 figure 1.24 and 1.25 for the 3 ways to represent
the formation of ions.
•
•
•
•
•
Formation of Ionic Compounds
Ions do not form by themselves.
As metallic and non-metallic atoms collide with one
another, their valence electrons interact.
The metal loses its valence electrons (becoming the
cation)and an adjacent non-metal gains them(the anion).
This is a transfer of electrons.
The two ions formed are opposite in charge and are
greatly attracted to each other, forming a very strong ionic
bond.
The rearrangement of electrons allows each ion a full
valence orbital (like its nearest noble gas) and leads to
greater stability and lower energy level for the ionic
compound.
• Ionic compounds:
• ions arrange themselves in a regular repeating pattern called a
crystal lattice
• Ionic compounds are usually hard, brittle solids at room
temperature that conduct electricity in solution.
• a binary ionic compound is formed from only 2
elements.
• Q? Why does Na become Na+ and Mg become
Mg2+?
• A. Sodium (Na) has only 1 valence electron, so when
it is lost, the resulting charge is 1+, however
magnesium (Mg) has 2 valence electrons so its ion has
a charge of 2+
Formation of an Ionic Compound
• Now a tougher one: The formation of
magnesium nitride
• Notice that the metal atoms donate six
electrons and the non-metal atoms accept six
electrons. The number of positive and negative
charges is balanced.
• Do BLM 1-7 “Isotopes and Ions” to review and
practice isotope and ion formation.
• Formation of Molecular Compounds
• When non-metals react with one another, electrons
are NOT transferred since both atoms tend to gain
electrons to fill their valence energy level.
• Instead, the valence electrons are shared. These
compounds are referred to as molecular and the
sharing of electrons forms a bond called covalent
(co = share, valent = valence shell)
• Molecular compounds:
• molecules do NOT usually form a crystal lattice shape
• Molecules can be solids, liquids or gases at room
temperature and usually do not conduct electricity in
solution.
• a diatomic molecule is composed of only 2 atoms and a
polyatomic molecule is composed of many atoms
• (see and memorize table 1.5 p 31)
•
It is important to note that the physical and
chemical properties of a compound is different from
those of the individual elements that make up the
compound.
Formation of a Molecular Compound
• Do Check Your Understanding p. 36 #1,3,6,7
• Do Chapter 1 Review p. 38
#2,4,5,7,10,11,12,15,16,17,19,20