Chemistry Practice Problems
Chemistry Midterm Study Guide
Steps of the Scientific Method:
- Observation
- Hypothesis
- Experiment
- Data
- Conclusion
- Retest
Experimental Variables:
- Independent Variable-the factor that is changed
- Dependent Variable-the factor that is measured or observed
- Control-the factor that is not being tested or changed. It’s used for comparison.
Mixtures:
homogenous mixtures
- have uniform composition
- components cannot be easily identified or separated
- example: sugar
water or vanilla ice cream
heterogeneous mixtures
- do not have the same appearance or composition throughout
- components can be easily identified and separated
- example: fruit
salad or chocolate chip cookies
Separating Mixtures
- two methods used in lab to separate mixtures
- filtration- use filter paper to filter out larger components
like sand
- distillation-evaporate liquid until the solid is left behind. Evaporate water and salt is left behind in flask.
Physical & Chemical Properties:
physical property
- characteristic that is observed with the senses and can be
determined without destroying the object.
- examples-color, shape, length, density, melting point
chemical property
- characteristic of matter that can only be observed when one
substance changes into a different substance, such as iron into rust.
- examples-flammability, reactivity with vinegar, reacts with
acid to form hydrogen gas.
Physical & Chemical Changes:
Physical Change
- change in which the identity of the substance does not change
- examples: density,
melting point, cutting paper, dissolving salt in water
Chemical Change
- transforms one type of matter into another kind, which may
have different properties
- examples: burning
leaves, bleaching your hair, frying an egg
Atomic History & Models:
Dalton’s Atomic Theory
- an element is composed of tiny, indivisible, indestructible
particles called atoms.
- all atoms of an element are identical and have the same properties.
- billiard ball model-atoms are indivisible spheres with constant
density throughout
Thompson
- plum pudding model-atoms pictured as spheres of positive charge
with small negative particles dispersed throughout the postive mass.
Rutherford
gold foil experiment-
- alpha particles were shot at a thin piece of gold foil.
- the particles were expected to pass through the foil, but instead
some particles were deflected and some bounced back.
- because some of the particles were deflected and bounced back,
this proved the existence of a dense nucleus in the center of the atom.
- since most of the particles passed right through, this indicated
that the atom is mostly empty space.
Bohr/Rutherford
planetary model-
- first developed by rutherford and later adapated by bohr.
- electrons orbit around the nucleus like planets around the
sun.
- the orbitals were described by rutherford as overlapping circles
surrounding the nucleus.
- the model was later changed by bohr, and he described the orbitals
as concentric circles that do not overlap and move out from the nucleus.
Atomic Structure:
atom-the smallest particle that represents an element
subatomic particles:
proton
- much larger and heavier than electrons
- postive electrical charge (+)
- located in the nucleus of the atom
neutron
- large and heavy like a proton
- no electrical charge (neutral)
- located in the nucleus of the atom
electron
- tiny, very light particles
- negative electrical charge (-)
- moves around the outside of the nucleus
atomic number
- describes the number of protons in an atom
- in a neutral atom the number of protons also equals the number of electrons
atomic mass number
- describes the number of protons + neutrons in an atom
isotopes-atoms of the same element that vary in the number of neutrons in the nucleus
calculating atomic mass
- convert %abundance into a decimal by moving the decimal point
over two places to the left.
- multiply the decimal by the mass of each isotope
- add the new masses together & round to two decimal places
- do not forget units of AMU
calculating PEN for
neutral atoms:
- the number of protons is equal
to the atomic number of the element
- the number of electrons equals
the number of protons if the atom is neutral
- the number of neutrons is equal
to (the atomic mass – the atomic number)
ions-charged particles that are formed when an atom loses or
gains electrons
cation-positively charged ion that results from the loss of electrons
anion-negatively charged ion that results from the gain of electrons
Calculating PEN For
Ions:
- each electron has a charge of (-1)
- the number of protons is equal
to the atomic number
- the number of electrons is equal
to the number of protons (+/-) any extra electrons to account for the charge
- the number of neutrons equals the
(atomic mass – atomic number)
EX:
- if an atom has a charge of (-3),
three electrons must be added to the normal electron count
- if an atom has a charge of (+2), two electrons
must be subtracted from the normal electron count
The Periodic Table:
most elements on the table are metals
elements are classified into three major groups
- metals
- metalloids
- found on the “step”
of the table between the metals and nonmetals
- have characteristics in between
those of metals and nonmetals
- nonmetals
periods
- horizontal rows on the periodic
table are called periods
- elements in a period are not alike
in properties
groups/families
- vertical columns on the periodic
table are called groups or families
- elements in the same group exhibit
similar chemical characteristics because they have the same number of valence electrons
valence electrons
- are electrons found in the outermost
shell of an atom
- these are the electrons that are
free to bond with other elements
- the number of valence electrons
is determined by the group number at the top of the periodic table
important groups of
the periodic table:
Alkali Metals
- group one on the table
- very reactive metals
- they react violently with water
- they are shiny and malleable
- form cations by donating one or
two electrons from their outer shell
Alkaline Earth Metals
- group two on the table
- reactive, but not as reactive as
the alkali metals
- always combine with nonmetals in
nature
- form cations by donating one electron
from their outer shell
Transition metals
- found in the center block of the
table
- reffered to as “D”
block
- form positive ions
- charge is indicated in parentheses
as roman numerals
Halogens
- group seven on the table
- very reactive and volatile nonmetals
- only need one electron to have
a full shell so they are likely to gain an electron and form anions
Noble Gases
- group eight on the table
- not very reactive
- have full electron shells so they
do not gain or lose electrons
- they do not form ions
Periodic Trends
- octet rule-atoms bond with other atoms until they have a full octet,
which is 8 electrons.
- atoms with fewer valence electrons (left side of table, groups 1
& 2) will lose electrons making them more positive, cations)
- atoms with more valence electrons (right side of table, groups 5-7)
will gain electrons making them more negative, anions)
Atomic Radius: size of the atom
- atoms increase in size as you move down a group
- atoms decrease in size as you move from left to right across the
table
- why? Going across the
table, the number of protons increases, making the electrical attraction between the nucleus and the electrons stronger and
that pulls in the electrons closer to the nucleus making the atomic radius smaller
Ionic Radius: size of an ion
- ions are charged atoms. a
cation is a positively charged ion and an anion is a negatively charged ion
- trend is the same as atomic radius:
- atoms increase in size as you move down a group
- atoms decrease in size as you move from left to right across the
table
- Cations are smaller than their neutral atom because they have lost
electrons, the remaining electrons are pulled in more by the strength of the protons in the nucleus.
- Anions are larger than their neutral atom because they have gained
electrons, so there is an increased repulsion between the nucleus and the orbital with the electrons.
Ionization Energy: energy required to remove an electron
from an atom
- ionization energy decreases as you move down a group because the
outer electrons are farther from the nucleus so they do not experience as much pull.
since they are not pulled with a lot of strength, they are easier to remove, thus requiring less energy.
- ionization energy increases as you move across from left to right
because elements to the right of the table are close to being full with 8 electrons.
since they are so close to full, they are not very likely to want to give up their electrons, thus requiring more energy.
Electronegativity:
- The ability of an atom to atrract electrons towards itself.
- electronegativity decreases as you move down a group because the
electrons in the outer orbital are father away from the nucleus and do not experience high attraction forces so additional
electrons would be difficult to attract. Also, elements to the left of the table
have extra electrons that they give away, so they would not need to atrract any additional electrons.
- electronegativity increases as you move across from left to right
because atoms to the right of the table are close to being full and need electrons, so they would need high electronegativity
in order to attract the electrons they need to be full and become stable.
Electron Configurations
- Groups 1 & 2 (s sublevels)
- Groups 13-18 (p sublevels)
- Groups 3-12 d block (d sublevels)
- Bottom Two (f sublevels)
- large number- row number
- letter-sublevel area
- superscript-how many spaces have you moved in the sublevel area
noble gas notation
- Find the noble gas closest to the element, that comes BEFORE
- Put the noble gas in brackets, continue notation from the next area
steps for writing an electron configuration:
- find the element
- start at the beginning of the periodic
table and indicate row number, letter section, and how many times you count across until you reach your element.
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