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Solution Manual For Foundations of College Chemistry, 16th Edition by Morris Hein, Susan Arena, Cary Willard Chapter 1-20
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Chemistry is the science that studies matter (everything that has a mass and occupies avolume). There are two transformations: physical and chemical, the first one it’s the change
of a physical property, the chemical composition of matter does not change, the second one
is the change in chemical composition. An element is a substance that cannot be divided
into other components through ordinary chemical reactions. A compound is a substance
formed by two or more atoms from the same or different properties from its constituents. A
mixture is made by two or more substances combined in not defined proportions. Properties
of constituents remain unaltered. Life of a cell, and of organisms, depends on the timely and
spontaneous occurrence of thousands of reactions.
According to Dalton’s atomic theory:
-Every element is made up of particles called atoms.
-All atoms of a given element are identical.
-Atoms of different elements have different properties.
-Chemical reactions do not change atoms of one element into those of another;(in chemical
reactions the atoms are neither created nor destroyed)
-Compounds are given by the combination of at least two atoms.
-In a given compound the relative amount of atoms and their type are constant.
ATOM: unit composed of a central core surrounded by electrons
- NUCLEUS: protons, endowed with a positive charge (+) +
neutrons, not charged
- ELECTRONS, endowed with a negative charge (-)
Protons and neutrons have roughly the same mass (1,6 x 10-24g), the mass of an electron
is around 1/1800 of the the mass of a proton.
The atomic radius is about 1 Å (10^-10 m). The radius of the nucleus is about 10^-15 m.
Atoms differ from each other in the number of elementary particles (protons, neutrons,
electrons) that compose them. Z = atomic number = n°of protons in the nucleus, A = mass
number = n°of protons + n°of neutrons. Atoms are electrically neutral if they contain the
same number of protons and electrons. The isotopes are atoms with the same atomic
number, but with different mass number; the most common and the most abundant isotope
is the protium, the second one is deuterium and the third one is the tritium (It is radioactive
and unstable because it has two neutrons).
Isotope STABILITY:
- To be stable a nucleus must have an optimal ratio between protons and neutrons;
- For nuclei with atomic number Z up to 20 the ratio neutrons/protons is normally 1;
- When Z increases, the number of neutrons necessary to keep the nucleus stable is larger
than Z.
Radioactive decay is the process through which unstable nuclei emit particles to transform
into more stable nuclei. When we have a stable isotope we can have different chemical
reactions: for example radioactive decay of tritium consists in the emission of an electron
called beta particle and happen when the nucleus is unstable for an excess of neutrons; the
tritium’s atom will transform into an helium atom (a neutron changed in a proton, and the
atomic number Z increases by a unit). Another kind of decay is alpha decay: is a type of
radioactive decay in which an atomic nucleus emits an alpha particle (helium nucleus) and
then transforms or "decays" into a different atomic nucleus, with a reduced mass number of
four and an atomic number reduced by two. Relative atomic mass: the Atomic Mass Unit
(AMU) is a unit of mass used to express atomic and molecular weights, equal to one twelfth
(un dodicesimo) of the mass of an atom of 12C. Molecular weight (g/mol) is the result of the
sum of the weights of the atoms that compose the molecule. The speed of light is measured
,like that: c=λ⋅v and the energy is: h (Planck constant)⋅v (frequency). A joule (J) is the work
required to exert a force of 1 newton for a distance of 1 meter. The joule is the unit of
energy, work and heat and is defined as 1 kg ⋅m^2/s^2=1N⋅m. N= is the amount
of force required to accelerate one kilogram mass one meter per second
squared kg⋅m/s^2. The calorie (cal) is a unit of energy. BOHR MODEL:
- The electron moves around the central proton following circular orbits;
- Not all the orbits are allowed but only some orbits related to the principal quantum number
(n); each of these orbits correspond to energetic level;
- The electrons are normally found by the lowest value of n and energy;
- If i want to promote an electron from the first orbit to the second i will have to give it at
least the energy that corresponds to the difference between the energies associated to the
two energy levels. If I give more energy, part of it will be lost as heat, because the electron
will use only the energy necessary to go to the other orbit; if I give less of what electrons
need, they will stay in the starting orbit, they will not be promoted to the next one;
- Electrons have a dual nature: they have mass but they have also waves; this dual nature
means that (Uncertainty principle of Heisenberg) we cannot know at the same time the
position and the velocity of the electrons —> if we know the velocity we can’t know the
position and vice versa. We never know where the electrons are, we can know the
probability of where we can find electrons. Electrons moving in its circular orbit around the
nucleus must be associated with a wavelength (λ) that depends on the mass of the particle
(m) and velocity (v), λ=h/mv (de broglie’s equation).
The wave function ψ (psi), describes the energy states accessible to electrons: the function
is known as the "Schrodinger equation”.
The allowed energy states coincide with those predicted by the Bohr model, however, for the
uncertainty principle, you can not specify both the positions and velocity of a particle. The
probability for the electron to be found in a given position around the nucleus is = ψ^2.
ORBITALS: The orbit is a defined circular path (percorso circolare definito) around the
nucleus in which the electrons rotate around the nucleus. The three-dimensional
space around the nucleus where the probability of finding an electron is
maximum is called an orbital (it represents the movement of an electron). The
Molecular orbital theory (MO) is a method for describing the electronic structure
of molecules using quantum mechanics. Each wave function which describes the
motion of an electron, corresponding to a given energy level, is called an ATOMIC
ORBITAL. Orbitals are described by the values of 4 quantum numbers (n, l, m,s).
Principal quantum number (n): it is a positive number (1,2,3..) and indicates the
size of the orbital and the relative distance from the nucleus, the greater the
value of n and the greater the energy level; 2) The secondary quantum number
(of angular momentum=l) is an integer number between 0 and n-1 and indicates
the shape of the orbital.The energy levels are given by the value of n, the
sublevels designate the shape of the orbital: l=0(s sublevel), l=1(p sublevel),
l=2(d sublevel), l=3(sublevel =f). 3) Magnetic quantum number (m): it
determines the orientation of orbital in space; -l ≤ m ≤ +l; 4) Spin quantum
number: describes the angular momentum (or spin) of an electron or other
particle; it defines the rotation of the electrons around the axis; they can rotate
clockwise or anti clockwise. +1⁄2 or -1⁄2 . All orbitals having the same value of n
(principal quantum number) have the same energy. How are electrons placed
into these orbitals?
, 1. the AUFBAU principles: is used to determine the electron configuration of an atom,
by adding electrons around the nucleus in the most stable energy level.
2. Principle of minimum energy: each electron occupies the available orbital with the
lowest energy level;
3. Pauli exclusion principle: in an atom it is not allowed to have two electrons which
possess identical quantum numbers. So an orbital can be occupied by two electrons
with different spin numbers.
4. Hund’s principle: if we have iso-energetic orbitals, we don't occupy the first one with
two electrons and then the other, but we start to occupy all of them with just one
electron with the same spin and then continue with the other.
PERIODIC TABLE:
- Along the horizontal line (periods): start with an element that has the outer electron
configuration of 1s electron and continue by increasing the atomic number. The external
electrons have the same quantum number n.
- Along the vertical line (groups) all elements have the same external electronic
configuration.
Mendeleev arranged the elements in order of increasing atomic, highlighting the periodic
nature of their chemical and physical properties (Periodic Table).
He predicted the existence of new elements, the usage of the atomic weight (and not the
number) made rationalization very difficult. IONIZATION ENERGY: the ionization energy is
the energy needed to remove an electron from an isolated atom (or ion) in the gas phase. I
can remove more than one atom so I can have the first ionization energy and the second
ionization energy, it is related to the attraction of the electron in the atom or ion; it increases
for each subsequent (= successivo) electron to be removed; it increases moving towards the
right along a period; it decreases going towards the bottom of a group because the external
elements are further away from the nucleus and suffer the screen effect of the lower levels.
The electrons are filling orbitals that are much more distant from the nucleus, so they will feel
less attraction from the positive nucleus. ELECTRON AFFINITY: It is the energy associated
with the process of addition of an electron to an atom or an ion. Neutral atom takes an
electron to become a negative ion and a positive ion takes an electron to become a neutral
atom. Halogens will easily get an electron to become negative ions because they will
complete their external electronic configuration in doing so; the addition of an electron for
most atoms results in a release of energy and the electron affinity is more negative toward
the right along a period and upwards within a group. ELECTRONEGATIVITY (Pauling):
Electronegativity, symbolized as χ, is the tendency for an atom, of a given chemical element,
to attract shared electrons (or electron density) when forming a chemical bond. The more an
atom is electronegative the more it will attract. If the difference in electronegativity of the two
atoms that are forming the bonds is high, the electrons will jump from one atom to the other.
A bond will be covalent or ionic just by considering the difference of electronegative between
the two atoms that form the bond: this is around 1.7-1.9, but, if the difference is lower, the
bond will be covalent (share atoms), if it is higher it will be ionic (one electron go from an
atom to another). POLARITY: Polarity is the distribution of electrical charge over the atoms
joined by the bond. Polarity depends on the electronegativity: if I form a bond with two
identical atoms, the maximum probability to find electrons will be in the middle. Every time I
have a difference of electronegativity higher than 0.4 I will have a covalent bond more or less
polar, depending on how much it is different.
, CHEMICAL BONDS: the chemical bond is the attractive force that holds the atoms in
molecules together; bond energy is the energy released to the environment when the bond
is formed or when we need to break the bond. It is expressed in KJ/mol or Kcal/mol;
If two atoms form chemical bonds they release energy so the system will have less energy
than the two isolated atoms; this is the reason why the bond is formed; on the other hand,
the energy that the two atoms need to break the bond, flows from the environment to the
bond that is broken. Interaction between atoms:1) Intramolecular= between atoms in the
molecule (ionic, covalent homopolar or covalent heteropolar); 2) Intermolecular= between
molecules (interactions between dipoles, hydrogen bonds) This interactions have different
energies: the bond form by two atoms inside the same molecule (the intramolecular
interaction) is stronger than the bond between two molecules (the intermolecular bond).
• Covalent bond: is the bond that is established between two atoms that share a pair of
electrons (bond pair). Two atoms can share 1, 2 or 3 pairs of e-, with single, double or triple
bond formation
Covalent bond can be: 1)Homopolar= a bond that is contracted by more atoms which have
more or less the same electronegativity. If the bond is formed by two identical atoms they
obviously have the same electronegativity. We can say also that a covalent bond is
homopolar when there is a small difference of electronegativity (between 0 and 0,4);2)
Polar= a bond where electrons are shared and where there is a difference in
electronegativity (between 0,4 and 1,9): the electrons involved are not equally shared, but
closer to the more electronegative atom.
• Ionic bond: is a bond between ions with opposite charge, the ionic bond is the electrostatic
force of attraction between a positively charged metal ion and a negatively charged non-
metal ion. Metals form positive ions because they lose electrons to become stable. In the
formation of the ionic bond we have to consider the electronegativity: in an ionic bond the
difference in electronegativity between the two atoms must be > 1.9; the ionic bond is not
directional: ionic bonds are formed when there is a transfer of electrons from one atom to
another. Since, a positive ion is attracted by a negation ion or vice versa. Due to this reason,
force of attraction occurs from both of the sides.
• Coordination or dative bonds: we have two electrons shared but they come both from the
same atom; one of the two atoms provides an orbital with a couple of electrons, the other will
provide an empty orbital (donor and acceptor). The two electrons came from the same atom
and once formed it has the same property of a normal covalent bond.
THE OCTET RULE: Each atom tends to acquire, lose or share electrons until it reaches the
external electronic configuration, consisting of eight electrons, equal to that of the nearest
noble gas (that already has the complete electronic configuration and is very stable thanks to
that). Elements with external configuration s and s2 can lose 1 or 2 electrons (form stable
cations). Elements with s2p5 external configuration tend to acquire 1 electron (form stable
anions). For energetic reasons, some orbitals are filled before others. The 4s orbital is filled
before the 3d orbital, because it has less energy. In the formation of cation it is important to
consider the ionization energy because the higher the EI, the more intense the attraction of
the electron inside the atom or ion. In the formation of anions it is important to consider the
electronic affinity. The elements of the third and subsequent period (with populated d
orbitals) form compounds in which the central atom can be surrounded by more than 8 e-.
There are three general exceptions to the octet rule.
- Molecules, such as NO, with an odd number of electrons, some of that are often referred to
as radicals and are highly reactive (the unpaired electron is very reactive);
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