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
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.
• 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
• The WHMIS or Workplace Hazardous Materials Information
System is a system of warning symbols designed to protect
people who use harmful substances at work.
• See figure 1.4 p8 for examples of WHMIS symbols.
***Read “Safety in the Chemistry Laboratory” p.x-xii
***Do the following pages in your note pack:
- WHMIS symbols
- Interpreting an MSDS
Lab quiz: _____________________
• Classifying Matter
• Matter is anything that has mass and occupies space.
– Mixtures
• can be mechanical (heterogeneous) where the separate
parts are visible ex: concrete or
• solutions (homogeneous) where the different parts are
not visible. Ex: 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.
• See figure 1.5 p10.
Matter
Pure Substance
Element
Compound
Mixture
Heterogeneous
(Mechanical)
Do Practice Problems #1-4 p10
Do Check Your Understanding #1-3 p11
Homogeneous
(Solution)
1.2 Developing Atomic Theories p.12-24
• 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
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
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.
Why model was rejected or
modified:
• Could not explain how
radioactive materials emit
alpha particles
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 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.
+
Rutherford’s “Solar System” Model
1871-1937
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.
Why it was rejected:
• 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
• 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.
Why is was rejected?
• at higher levels there are sublevels
• 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
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
*This is what we use in High School
• Quantum Mechanics Theory: Most current accepted
atomic theory.
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+
Did you know?
• 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?
TRUE OR FALSE?
• The Russian scientist Demitri Mendeleev was
responsible for organizing the modern Periodic
Table.
• Metals are found on the right side of the Periodic
Table.
• All elements were known when the Periodic Table
was developed.
TRUE OR FALSE?
• The Russian scientist Demitri Mendeleev was
responsible for organizing the modern Periodic
Table. True
• Metals are found on the right side of the Periodic
Table. False
• All elements were known when the Periodic Table
was developed. False
TRUE OR FALSE?
• All elements are pure substances.
• There are many patterns in the Periodic Table.
• Elements are organized according to
increasing size.
TRUE OR FALSE?
• All elements are pure substances. True
• There are many patterns in the Periodic Table.
True
• Elements are organized according to
increasing size. False
TRUE OR FALSE?
• Elements are organized according to
increasing mass.
• The staircase locates all of the metalloids.
• Molar mass values describe the weight of 1
gram of that element.
TRUE OR FALSE?
• Elements are organized according to
increasing mass. True
• The staircase locates all of the metalloids.
False? (All??)
• Molar mass values describe the weight of 1
gram of that element. False
The Periodic Table tells us Many Things About
the Elements
11
23
Na
sodium
•
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 = number of neutrons
For example, the most abundant isotope of
uranium is uranium-238. The complete symbol
of this isotope is:
• How many protons are in uranium-238 ?
• 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!
29
• Name this isotope:
14
Si
29
• Name this isotope:
14
• silicon-29
Si
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!
• Do Practice Problems p. 23 #5-8
• Do Check Your Understanding p. 24 #1-5
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
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 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.
Based on these patterns, color in your periodic table on page 11 in
your note pack
Periods
• 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
Reactions with Water
Brainiak Alkali Metal Video
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.
• 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
• 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
2e8e2e12p+
Mg
• The diagram representing the element beryllium looks
like this:
• The diagram representing the element beryllium looks
like this:
2e2e4p+
Be
• Try the diagram for fluorine:
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
• Lewis structures are used 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:
• 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 and an adjacent
non-metal gains them. 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.
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.
Examples:
• The formation of sodium chloride and
magnesium oxide
• 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
• Do Investigation 1-A p 33 “Ionic or Covalent:
Track Those Electrons” + Questions 1-6 p. 35
• 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