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1
Asst. Prof. Dr. Abbas Ali Salih Al-Hamdani
Inorganic Chemistry 2nd year
1- Modern Periodic Table
-Electronic Configuration of atoms.
-Identifying isotopes:
-Ionic Energy.
-Electronic Affinity.
-Elector negativity.
-Shielding.
-Atomic Radii.
2- Acid- Base & Salts Chemistry.
Acid- base Concepts.
Arrhenius
Bronsted-Lowery definition.
Lux-Flood definition.
Solvent system definition.
Lewis definition.
Usanovich definition.
Ageneralized acid-base concept
Acidity of oxy acid.
Basically and hydrolysis reactions.
Hydration and Hydrolysis reaction.
Basicity of Substituted amines
2
Hard and Soft acid & base
Classification of acids and bases as hard or soft.
3- Chemical Bonds (Covalent Bonding)
The Lewis octal theory.
The Isoelectronic principle.
Bond Distance.
Bond strength.
Bond energies.
Bond polarity.
Dipole Moments.
Valence-shell Electron pair Repulsion.
Electro negativity and hardness and softness.
4- Hybridization.
5- Valence Bond Theory.
6- Molecular Orbital Theory.
7- Symmetry.
8- Symmetry element
Symmetry operation.
Point group symmetry
Identity, E
Proper rotation, Cn
Proper axis
Rotation axis.
3
Reflection,
Mirror plane.
Reflection plane.
Inversion, i
Inversion center
Center of symmetry
Rotation- reflection Improper rotation, Sn
Improper axis.
Alternating axis.
9- Hydrogen.
Physical, Chemical Properties and Reaction.
10- Main group elements (I):
Physical, Chemical Properties and Reaction.
11- Main group elements (II):
Physical, Chemical Properties and Reaction.
4
:المصادر المفيدة حول الموضوع
• Lectures Note (part one)- handout
• - Advance Inorganic Chemistry Volume(I,II) (2010), Sataya prakash, G.D.Tuli
• - Concise Inorganic Chemistry, by: J0D.Lee (2003).
• - Inorganic Chemistry (Principles of Structure and Reactivity) , 4Ed ,james, E.
Huheey (1983)
أحسان عبد الغني، جامعة الموصل .سعد عزالدين و د. الكيمياء اللاعضوية والتناسقية ، د •
.8811
. منذر الجنابي و د. نعمان النعيمي و د. د( : والقسم الثاني -القسم الاول)الكيمياء اللاعضوية • .6791. بغداد نظير ميلاد ، جامعة
5
1-Modern Periodic Table:
The Periodic Table: In the 1870's, the Russian chemist Mendeleev developed
the periodic table, based upon the relationship between the atomic weights of the
elements and their chemical properties. As one ascends from lightest to heaviest
elements, there is a periodic recurrence of chemical properties. For example, the
elements with atomic numbers 2, 10,18, 36, 54, and 86 all are chemically inert (the
noble gases), while those with atomic numbers one greater 3, 11, 19, 37, 55, and 87
are all extremely reactive metals (the alkali metals). The rows or periods of the
periodic table contain the elements with the same principle quantum number (n) for
their outermost electrons. That is, until the fourth period, which contains the first row
of transition metals, where the electrons of highest energy exist in d orbitals, which
follow in the building-up order s orbitals of a higher shell, being grouped in the same
period with elements of higher principle quantum number (n). The columns or
groups in the periodic table contain elements with the same electron configuration
in the outermost shell. For example, the outermost shell of a noble gas is completely
filled, while the outermost shell of an alkali metal contains only one electron.
6
7
- Has five key pieces of formation about any nuclide. Starting with the number
in the lower left and going clockwise around the S
✓Atomic number (lower left, 16): This is the number of protons within the nuclide.
As the atomic number tells you the element, which is also represented by the
element’s symbol at the center of the nuclear notation (in
8
this case, S for sulfur), it’s commonly not included. All atoms that have 16 protons are
sulfur atoms.
✓Mass number (upper-left, 32): This represents the sum of the neutrons and
protons present in the nuclide. To figure out the number of neutrons a nuclide has,
subtract the atomic number (16) from the mass number (32).
✓Ionization state (upper-right, 2+): Remember that an ion is an atom that has
gained or lost an electron, so it has a net electrical charge. If positive (+), the atom lost
electrons; if negative (–), the atom gained electrons.
✓Atoms per molecule (lower-right, 2): This number simply tells you how many
atoms are making up the molecule or polyatomic ion in question, just as the “2” in
H2O tells you there’s two hydrogen atoms in each water molecule.
✓Element’s symbol (center, S): Just like the atomic number (and also straight off the
periodic table), this tells you what element your nuclide
- Electronic Configuration of atoms:
بداية تعلم الترتيب الالكتروني وفق ترتيب الجدول الدوري ضع : الترتيب الالكتروني
الجدول الدوري اماك وأبد
Examples: 1H: 1S1 3Li: 1S2 2S1
11Na: 1S2 2S22P63S1
19K: 1S2 2S22P63S23P64S1
37Rh: 1S2 2S22P63S23P64S23d104P65S1
55Cs: 1S2 2S22P63S23P64S23d104P65S24d105P66S1
9
87Fr: 1S2 2S22P63S23P64S23d104P65S24d105P6 6S25d104f146P6 7S1
متشابه للمجموعة نفسها الاخير نجد ان الترتيب لاحظ التريب لجميع عناصر المجموعة
فقط يختلف رقم المدار
وممكن ان نكتب الترتيب كما يلي للسهولة
Examples: 1H: 1S1 3Li: [He]2 2S1
11Na: [Ne]10 3S1
19K: [Ar]18 4S1
37Rh: [Kr]36 5S1
55Cs: [Xe]54 6S1
87Fr: [Rn]86 7S1
Examples: (valence electrons in blue)
P: [Ne]3s23p3
As: [Ar] 4s23d104p3
I: [Kr]5s24d105p5
Ta: [Kr]6s24f145d3
Zn: [Ar]4s23d10
حسب مجموعته فقط أحفظ الترتيب الكتروني وهكذا بالنسبة الى بقية العناصر وكلا
مع وجود بعض الملاحظات نأتي عليها عند لنعصر المجموعة الاول والبقية تأتي تباعا
.الشرح
10
Identifying isotopes:
There are three naturally occurring isotopes of oxygen: oxygen-16, oxygen-17,
and oxygen-18. All three of these isotopes have an atomic number of 8, so these
isotopes have eight, nine, and ten neutrons, respectively. Figure 2-6 displays these
three isotopes of oxygen. Over 99.7 percent of all the oxygen are oxygen-16 due to
the special stability that comes from having an equal number of protons and
neutrons for smaller atoms. However as atoms get larger, they require more neutrons
to remain stable. Therefore, small oxygen-16 (8 protons, 8 neutrons) and large lead-
204 (82 protons, 122 neutrons) are both stable. Isotopes all act the same chemically,
because they have the same number of protons, which implies they have the same
number of electrons, which deter mines chemical reactivity. Isotopes that are
unstable undergo radioactive decay; they are called radioisotopes
11
- Ionic Ionization Energy:
The ionization energy is the energy necessary to remove an electron from an
atom: The higher the ionization energy, the stronger the electrons are bound to an
atom. As you move left to right across a period of the periodic table, the first
ionization energy generally increases due to the decreasing radius of the atoms.
Think of pulling on a really large magnet near your refrigerator it’s easy to move if
you are a few feet away, but takes a lot of strength to rip away when only inches
apart. The same is true for the electrons opposite charges attract (negative electrons
to the positive protons of the nucleus). Going down a group of the periodic table, the
atoms get larger, so the ionization energies decrease from top to bottom. Each
electron in an atom has anionization energy, so the energy required to remove the
first electron is known as the first ionization energy, the next electron is known as the
second ionization energy, and so on. After removing one electron, the electron-
electron repulsion is smaller in the electron cloud surrounding the nucleus, so the
electron cloud contracts a little making the ion a little smaller. As stated before, the
12
smaller the atom, the higher the ionization energy. Following this trend, each
subsequent ionization energy is higher in energy. If you monitor these energies, not
only do they increase at every new ionization energy, but eventually you see a very
large jump; for example, the first four ionization energies for aluminum are: 578
kJ/mol, 1,817 kJ/mol, 2,745 kJ/mol, and 11,577 kJ/mol, respectively. The massive
jump in energy occurs between the third and fourth ionization energies. Aluminum,
having only three valence electrons, is electronically like a noble gas after removing
three electrons, so removing the fourth electron (the fourth ionization energy)
requires a large amount of energy because you are destroying the very stable noble
gas configuration.
لة وهي في ادنى حالات الطاقة هذا لنزع الكترون من ذرة غازية متعادوهي أقل طاقة لازمة
ني فهو الطاقة اللازمةاجهد التاين الاول اما الث ىبالنسبة اللنزع أضعف الكترون ارتباطا
.من الايون الموجب الاحادي في ادنى حالات طاقته
وان طاقة وان لكل عنصر عدد من جهود التاين يساوبي عدد الكترونات هذا العنصر
لذالك هي موجبة الاشارة (……>I1 <I2 <I3): التاين تزداد وهي من النوع الماص للحرارة دائما
يعرف جهد التأين بأنه أقل طاقة لازمة لأنتزاع ألكترون من ذرة غازية وهي في أدنى حالات
.الطاقة لأنتاج أيون موجب
(IP1)تسمى طاقة التأين اللازمة لأنتزاع ألكترون واحد من الذرة الغازية بطاقة التأين الأولى
.وهكذا (IP2)وتدعى طاقة التأين اللازمة لأنتزاع الألكترون الثاني بطاقة التأين الثانية
13
:لقد لوحظ أن طاقة التأين تزداد تبعا للترتيب الآتي
:وتعتمد قيمة طاقة التأين على العوامل الآتية
.مقدار الشحنة النووية المؤثرة والتي تعتمد بدورها على مدى حجب الألكترونات الأخرى -1
المسافة بين الألكترون والنواة أو بمعنى آخر أدق طول نصف القطر الأكثر أحتمالا لهذا -2
.الألكترون
مدى نفاذية الألكترون للسحابة الألكترونية للألكترونات الأخرى حيث أن نفاذية -3
S P d f: الألكترونات تقل في هذا الأتجاه
.حيث تزداد طاقة التأين بزيادة شحنة الأيون الموجب, مقدار شحنة الأيون الموجب -4
ويعطي الشكل الآتي أيضاحا مبسطا على أن جهد التأين يقل في الزمرة الواحدة مع زيادة
العدد الذري في حين يزداد جهد التأين في الدورة الواحدة كلما زاد العدد الذري نتيجة
لأنكماش حجوم الذرات في الدورة الواحدة نتيجة الزيادة في الشحنة النووية المؤثرة
Amount of energy required to remove an electron from the ground state of a
gaseous atom or ion.
– First ionization energy is that energy required to remove first electron.
14
– Second ionization energy is that energy required to remove second electron,
etc.
• It requires more energy to remove each successive electron.
When all valence electrons have been removed, the ionization energy takes a
quantum leap.
15
-Electron affinities
Opposite to ionization energy, electron affinity is the amount of energy an
atom or ion releases when an electron is added. When an atom takes on an electron,
general energy is released, which would be a negative change in energy (ΔΕ), so
electron affinities are usually negative. To complicate matters even further, for many
atoms it is easier to determine the ionization energy (energy to remove an electron)
for the atom with an extra electron. In other words, if electron affinity is, then you
could do the reverse. By doing the opposite, the change in energy would be positive,
not negative. In general, nonmetals have higher magnitudes for electron affinities
than metals
وهي في ادنى حالات الطاقة بالكترون وهي الطاقة المتحررة عند اتحاد ذرة غازية متعادلة
.معطية الايون السالب الاحادي الغازس في ادنى حالات الطاقة أي انها طاقة التفاعل
Electro negativity:
If you take the average of an atom’s ionization energy and electron affinity,
you have a measure for the likelihood of an atom attracting an electron, known as
electronegativity. The higher the electronegativity, the more an atom would like an
electron. This mathematical method of calculating electro negativities is known as
Mulliken electro negativity, named after its Nobel Prize–winning discoverer, Robert
Mulliken. Another Nobel Prize–winning scientist, Linus Pauling, came up with a
similar method that was calculated using the relative strengths of bonds in
molecules; for example, if you knew the bond energy for hydrogen (H-H), chlorine
(Cl-Cl), and hydrogen chloride (H-Cl), then you could determine which element was
more likely to attract an electron relative to one another. If you continue this
calculation through the entire periodic table, you would have the Pauling electro
16
negativity scale (which ranges from 0.7 for the least electronegative francium to 4.0
for the most electronegative fluorine). Regardless of which scale you use, (there are
numerous others not even mentioned, though Pauling and Mulliken are the most
commonly encountered) the general trend is the same: Electro negativity increases
going from the bottom to the top or from the left to the right on the periodic table.
The general trend is shown in Figure
وهذه الميزة تتميز بها الذرة فقط عندما ) هي قوة ذرة في جزيئة ما على جذب الكترون نحوها
(.تكون متحدة مع غيرها من الذرات وليس في حالتها المنفردة
Shieldingالحجب
تؤثر مقدار قوة جذب النواة لأي ألكترون متواجد ضمن الأغلفة الثانوية على مقدار
يحس بالشحنة 1Sفمثلا الألكترون المتواجد في , الطاقة اللازمة لأنتزاع ذلك الألكترون
أكثر من الألكترون المتواجد ( التي تمثل قوة جذب النواة لذلك الألكترون ) *Zالنووية المؤثرة
:من خلال الصيغة *Zويمكن حساب الشحنة النووية المؤثرة 2Sف في الغلا
Z*= Z- S
17
Z* : الشحنة النووية المؤثرة
Z : العدد الذريS : ثابت الحجب
بإتباع قواعد معينة وضعت من قبل العالم Shielding Constantيمكن حساب ثابت الحجب
:وكما يأتي Slaterسليتر
:nPأو nSلحساب ثابت الحجب لأي ألكترون يقع في الغلاف من النوع قواعد سليتر ●
:يكتب الترتيب الألكتروني للعنصر حسب الترتيب الاتي -1
(1S) (2S 2P) (3S 3P) (3d) (4S 4P) (4d) (4f) ……..,ect.
.لايدخل ضمن حساب ثابت الحجب nPأو nSأي ألكترون في يمين -2
كل الألكترونات ضممن المدار ) كل الألكترونات الواقعة ضمن نفس المدار للألكترون المعني -3
n ) ؤثرة من الشحنة النووية الم( 5.30)تحجب بمقدار.
( .0..5)للألكترون تحجب بمقدار (n-1)كل الألكترونات ضمن المدار -4
(.1)فما دون تحجب بمقدار (n-2)كل الألكترونات ضمن المدار -0
:7Nأحسب الشحنة النووية المؤثرة لألكترون التكافؤ الأخير لذرة : مثال
7N 1S2 2S2P3
(1S)2 (2S 2P)5
S=( 4*0.35 ) + (2* 0.85)= 3.10
Z* =Z-S = 7-3.10 = 3.9
:7Nأحسب مقدار الشحن النووية المؤثرة للألكترون الرابع في ذرة : مثال
7N 1S2 2S2P3
(1S)2 (2S 2P)5
(2S)ع يعني الموجود في وللألكترون الراب
S=(1* 0.35) + (2* 0.85) =
Z*= Z-S= 7-
:15Pون الأخير في ذرة الفسفور للألكتر *Zأحسب مقدار : مثال
15P 1S2 2S2 2P6 3S2 3P3
18
(1S)2 (2S2P)8 (3S3P)5
S= (4*0.35) + (8*0.85) +(2*1) = 10.20
Z*= Z-S = 15 – 10.20 = 4.80
والألكترون (4S)ؤي التي يحس بها الألكترون التكاف *Zأحسب الشحنة النووية المؤثرة: مثال
30Znلذرة الخارصين (3d)الأخير في
30 Zn: 1S2 2S2 2P6 3S2 3P6 3d10 4S2
(1S)2 (2S 2P)8 (3S 3P)8 (3d)10 (4S)2
: 4Sللألكترون الأخير *
S= ( 1* 0.35) + (18*0.85) +( 10*1) = 25.65
Z*= 30-25.65 = 4.35
:3d للألكترون الأخير *
S= (9* 0.35) + (18*1) =21.15
Z* = Z-S = 30 – 21.15 = 8.85
الدورة الواحدة بزيادة العدد الذري وتقل في ولهذا فإن مقدار الشحنة النووية المؤثرة يزداد في
.الزمرة الواحدة بزيادة العدد الذري
Atomic Radiiنصف القطر الذري
وجد أن الكثير من الخصائص الفيزيائية كالكثافة ودرجة الغليان ودرجة الأنصهار لها علاقة
فقد تم قياس نصف بحجوم الذرات وبما أنه من غير الممكن عزل الذرات كذرات مفردة لذا
القطر الذري بصورة غير مباشرة من خلال قياس المسافة بين نواتي ذرتين متآصرتين في
(.البلورة) جزيئات الغازات أو بين نواتي أيونين في الحالة الصلبة
19
لوحظ أن نصف القطر الذري يزداد بزيادة العدد الذري لعناصر المجموعة الواحدة كلما
ألى الأسفل بسبب زيادة عدد الأغلة الثانوية في حين يقل نصف القطر الذري أتجهنا من الأعلى
في الزمرة الواحدة بزيادة العدد الذري أي كلما أتجهنا من اليسار ألى اليمين بسبب الزيادة في
.الشحنة النووية المؤثرة في الوقت الذي يبقى فية عدد الأغلفة الثانوية ثابتا
20
2- The ABCs: Acid-Base Chemistry Acid / Base Reactions:-
Acids and bases have been known for thousands of years due to their practical
applications. To the ancient Greeks acids were referred to as oxein, meaning “sour
tasting.” In Latin they were known as acere, which came from the word acetum,
meaning vinegar. Later, this was adapted to the English vocabulary as acid.
As these descriptive names suggest, acids were originally distinguished by
their sharp, sour taste. Not only are they sour, but they were found to be corrosive to
metals. Similar to rusting, acids oxidize metals. They were then found to react with
lichen plants and caused litmus paper to change color. Bases were often investigated
and categorized according to how they neutralize, change, or counteract acid
solutions. They were first recognized when making soaps using roasted ash. The
word alkaline hails from the Arabic word “to roast.” The soap-making process creates
alkaline-water-based mix-tures composed of both slaked lime and roasted ash.
Consider the expression “every reaction has an equal and opposite reaction.”
Consider acids and bases are each equal and opposite to each other; when you add
together an acid and a base, they cancel each other out, and in that respect they are
opposite to each other.
Acid / Base definitions
Definition 1: Arrhenius
In 1884 Arrhenius proposed his theory of electrolytic dissociation based on the
self-ionization of water:
Arrhenius acid is a substance that produces H+ (H3O+) in water
21
Arrhenius base is a substance that produces OH- in water
Brønsted-Lowry theory
In 1923, two independent scientists, Johannes Brønsted and Thomas Lowry,
defined acids to be proton donors and bases to be proton acceptors. Continuing to
work with the concept of hydrogen in acids and bases, Brønsted and Lowry each
independently defined a new set of conditions that has come to be known as the
Brønsted-Lowry model. They continued with the notion of acids being proton
donors, but they expanded on the concept of the base to include any substance that
is capable of binding protons. In other words, any compound that complex with a
Brønsted acid is considered a Brønsted base. In aqueous solutions this is not a major
distinction, but it was known that other acid base reactions can occur in solvents that
are not water. The Arrhenius theory applies to water. But as with the water example,
where the self-ionization of water occurs, hydronium ion concentration increases
above a value of 10-7 M for acids, and those that decrease the value are bases.
The Brønsted-Lowry theory expands on this notion beyond solvents of just
water, to cover others such as liquid ammonia, glacial acetic acid, anhydrous sulfuric
22
acid, and all solvents containing hydrogen. Bases can accept protons but don’t
necessarily contain OH-.
For example,
A Brønsted-Lowry acid is a proton donor, A Brønsted-Lowry base is a proton acceptor
In liquid ammonia, the NH4
+ donates a proton, whereas the NH2- accepts a proton:
When the chemical species differ in composition by a proton, they are called a
conjugate pair. For every acid there is a conjugate base, and vice versa
23
In this reaction, the ammonium salt acts as the acid, because it donates the
protons, whereas the sulfide ion is a base because it accepts the protons. This
reaction is reversible and proceeds to produce the species that is weakest, such as HS-
and NH3.
This explanation of acids and bases depends on the choice of solvent. If a
solvent is chosen that is a strong donator of protons, then it is considered the acid in
the reaction, even if it’s not technically called an acid. Or, if two acids are mixed
together, the weaker donor of protons is instead classified as the base. Take, for
example, the case of hydrogen fluoride mixed in perchloric acid
HClO4 + HF → H2F + ClO4-
Both of these materials can be individually considered acids, but when they
are mixed together the hydrogen fluoride gains a proton and the perchloric acid gave
one up. In this case, the perchloric acid was a stronger acid because it was the acid
that donated the H+, and the HF was a weaker acid, so it acted like a base and
accepted the H+
24
Accepting or donating: Lewis’s theory
Up to this time, the ideas about acids and bases focused on the use of water as
a solvent. Water is a protic solvent because it has the capacity to donate protons.
Water is polar, which means it has a slight electric charge (see Chapter 10 for more
details). Gilbert Lewis considered the case of a nonpolar, or nonprotic, solvent and
how dissolution might occur in non-aqueous solvents where polar species cannot be
dissolved.
Lewis basically threw out the two previous ideas and instead looked at the
reactions based on the transfer of electron pairs. He postulated that acids accept
electron pairs, and bases donate electron pairs.
In the Lewis model, an acid accepts electron pairs, whereas a base donates
electron pairs. Any proton is a Lewis acid, whereas ammonia, for example, is a Lewis
base, because the lone pairs of the nitrogen are donated to a proton as shown in the
following equation. (In this example the backward pointing arrow is used to signify
that the electrons are being donated.)
Organic chemists commonly refer to Lewis acids as electrophiles and Lewis bases as
nucleophiles. A classic example of a Lewis acid-base reaction involves BF3 and NH3.
The ammonia has a lone pair that can donate to the empty p orbital of the BF3. In this
reaction, the BF3 is the Lewis acid and NH3 is the Lewis base. This results in the
formation of a covalent bond between them.
25
NH3 +H2O NH4+ + OH-
Base acid acid base
Comparing Lewis and Brønsted theories
Because you can look at acids as being either proton donors or as electron
acceptors (and bases vice versa), you commonly see a complex labeled as either a
Lewis acid/base, or Brønsted acid/base. For this reason it’s valuable to consider the
similarities and differences between these definitions. Here are the key points to
remember:
26
Usanovich Definitionتعريف نظرية
القواعد أو تمنح القواعد أو تمنح أيونات موجبة أو تستقبل الحامض هو أية مادة كيميائية تتفاعل مع
.أيونات سالبة او الكترونات
.القاعدة على عكسها
27
فهو يشمل جميع تفاعلات لويس بين الحوامض والقواعد اضافة الى تفاعلات الاكسدة والاختزال التي
لالكترونات او اكثر وركزي على عدم الاشباع الذ كاملا
ي يميز بعض التفاعلات للحوامض تشمل انتقالا
والقواعد
OH- + O=C=O → HOCO2
Lux-Floofتعريف
تصف تصرف الحوامض والقواعد بدلالة ايون الاوكسيد اي القاعدة تعطي للاوكسيد والحامض
تنحصر الفائدة في الانظمة للاكاسيد المنصهرة وهو يركز على الصفات الحامضية , .مستقبل للاوكسيد
..للانهدريدات والقاعدية
Ca2+ + O2- + H2O → Ca 2+ + 2OH-
القاعدة هي انهدريد حامض ي
SiO2 + H2O → H2SO4
الحامض انهدريد حامض ي لكن الاتجاه العكس ي لسحب الماء هو الارجح
Solvent System Definition:
وتكون أيونات سالبة وايونات موجبة تماما كما .يتأين الماءيتأين كثير من المذيبات ذاتيا
2H2O H3O+ + OH-
2NH3 NH4+ + NH-
2
2H2SO4 H3SO4+ + HSO4
-
2OPCl3 OPCl2+ + OPCl4
-
الحامض هو المركب الذي يزيد من تركز الايون الموجب للميز للمذيب و القاعدة هي المركب
الذي يزيد من تركز الايون السالب للميز للمذيب
Generalized Acid Base Concept:
او انه يستقبل ايونات سالبة , والقاعدة تمنح ايونات سالبة او يمنح اونات موجبة : الحامض
تستقبل ايونات موجبة
الحامضية هي الصفة الموجبة للمادة الكيميائية وانها تنقص بتفاعلها مع القاعدة, القاعدية
بتفاعلها مع القاعدة هي الصفة السالبة للمادة الكيميائية وانها تنقص
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Pearson's Hard/Soft Acid/Base Theory
This can be used to predict reactivity.
Hard acids or bases: high charge density, not polarizable Soft acids or bases: low
charge density, polarizable Operating Principle: hard acids prefer to bind with hard
bases and soft acids prefer to bind with soft bases
HSAB can be used to qualitatively predict reactivity, especially metathesis reactions:
CuF (s) + HI (aq) → CuI (s) + HF(aq)
S-H H-S S-S H-H
CaO (s) + 2 HBr (aq) → Ca2+(aq) + 2 Br–(aq) + H2O(l)
H-H H-S H S H-H
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Chemical Bonds:
Chemical bonds hold molecules together and create temporary
connections that are essential to life. Types of chemical bonds including
covalent, ionic, and hydrogen bonds and London dispersion forces.
والتي أخرى ذرات أو التكافؤ لذرة وإلكترونات النواة تجاذب بين هي قوى الكيميائية الرابطة
البعض ببعضها مرتبطة الذرات تلك تجعل
لعدة تبعا تختلف وإنما الطريقة, بنفس كلها تتم لا الكيميائية الروابط الكيميائية الروابط أنواع
تقسيم يمكن العامل هذا أساس وعلى. للذرة الخارجي الغلاف في الإلكترونات عدد أهمها عوامل
أيونية ورابطة تساهمية رابطة إلى الروابط أنواع
lonic bondالرابطة الأيونية
جزيء ليكونوا سالبة وأيونات موجبة أيونات بين اتحاد فيها يحدث الأيونية الرابطة
Na الصوديوم NaCl الطعام ملح هو الأيونية الرابطة على مثال الكهربائية الشحنة متعادل
Bonds
Chemical bonds
Inonic coordination
Metallurgical Covalent
polar
non polar
Physical bonds
Hydrogen bonds
Vander val
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17يحمل Cl17 الكلور بينما )كالتالي الإلكترونية المدارات في مرتبين إلكترونا على يحتوي 11
.)كالتالي المدارات في مرتبين إلكترونا
1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance
in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or
ions
Covalent Bondالاصرة التساهمية
• Between nonmetallic elements of similar electronegativity.
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• Formed by sharing electron pairs
• Stable non-ionizing particles, they are not conductors at any state
• Examples; O2, CO2, C2H6, H2O, SiC
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Polar bonds and Electro negativity
Bonds in all the polyatomic ions and diatomics are all covalent
bonds
Nonpolar covalent bonds
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when electrons are shared equally
Oxygen Molecule (O2)
POLAR COVALENT BONDS
when electrons are shared but shared unequally
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water is a polar molecule because oxygen is more electronegative than
hydrogen, and therefore electrons are pulled closer to oxygen
METALLIC BOND
bond found in metals; holds metal atoms together very strongly
• Formed between atoms of metallic elements
• Electron cloud around atoms
• Good conductors at all states, lustrous, very high melting points>
• Examples; Na, Fe, Al, Au, Co
• Compute the mass of the following compounds round to nearest
tenth & state type of bond:
• NaCl;
• 23 + 35 = 58; Ionic Bond
• C2H6;
• 24 + 6 = 30; Covalent Bond
• Na(CO3)2;
• 23 + 2(12 + 3x16) = 123; Ionic & Covalent
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The three models of chemical bonding
Molecular Geometry Molecules of different substances have diverse shapes. Atoms
attach to one another in various geometric arrangements. The overall
molecular shape of a molecule is determined by its bond angles in three dimensions. The shape of a molecule is very important for its physical
and chemical properties. Molekül geometrisini açıklamak için farklı
kuramlar mevcuttur. modeller
VSEPR Valence Bond Theory Molecular Orbital Theory
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Valence Shell Electron Pair Repulsion (VSEPR) Theory
It is a method for predicting the shape of a molecule from the knowledge of the groups of electrons around a central atom.
A method to predict the shapes of molecules from their electronic
structures (Lewis structures do not depict shape)
Basic principle: each group of valence electrons around a central
atom is located as far away as possible from the others in order to minimize repulsions Both bonding and non-bonding valence electrons around
the central atom are considered. AXmEn symbolism: A = central atom, X = surrounding atoms,
E = non-bonding electrons (usually a lone pair)
Electron pairs (bonding and nonbonding electrons) repel one another, as
a result, the electron pairs remain as far apart as possible from another as possible to minimize the repulsion.
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• Two electron pairs in the valence orbital are arranged linearly
• Three electron pairs are organized in a trigonal planar arrangement
• Four electron pairs are organized in a tetrahedral arrangement
• Five electron pairs are arranged in a trigonal bipyramid • Six electron pairs are organized in an octahedral arrangement
The repulsion of lone pair electrons is grater than the repulsion of
bond pair electrons
Electron pairs assume orientations about an atom to minimize repulsions.
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Applying the VSEPR theory.
• draw a plausible Lewis structure of the molecule or polyatomic
ion. • Determine the number of unshared electron pairs and numbers of
bonds around the central atom ( multiple bonds count as a single
bond)
• B + No
• Establish the geometrical orientation of the electron pairs around
the central atom as linear, trigonal planar, tetrahedral,, trigonal bipyramid or octahedral
• Describe the molecular geometry
question: predict the shape of CO2 HCN CH4, NH3 SO2 PCl5, SF6 and
H2O by using VSEPR theory.
question: predict the shape of CO2 HCN CH4, NH3 SO2 PCl5, SF6 and
H2O by using VSEPR theory.
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The nonbonding electron pairs are as important as bonding electron pairs in determining the structure.
Nonbonding electrons take up more space in the valence shell than the
bonding electrons. If one or more of the electron pairs are lone pairs, the distribution of
electron pair and the geometrical shape of the molecule must be
different. The bond angles decrease as the number of nonbonding electron pairs
increases
Repulsion strengths
lone pair -lone pair lone pair e-bond pair bond pair-bond pair
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The three molecular shapes of the tetrahedral electron-group arrangement
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The four molecular shapes of the trigonal bipyramidal electron-
group arrangement
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Valence bond theory
The covalent bonds are formed by overlap of atomic orbitals each of
which contains one electron of opposite spin.
Basic Principle A covalent bond forms when the orbitals of two atoms overlap and are
occupied by a pair of electrons that have the highest probability of being
located between the nuclei.
Three Central Themes A set of overlapping orbitals has a maximum of two electrons that must
have opposite spins. The greater the orbital overlap, the stronger (more stable) the bond.
The valence atomic orbitals in a molecule are different from those in
isolated atoms (hybridization).
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The valence bond method predicts molecule shapes from the shapes and
orientation of the atomic orbitals and their overlap regions when two
atoms approach.
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In most cases the orbitals that overlap are reconfigured orbitals, called
hybrid orbitals, having different shapes and orientations than pure orbitals.
The process of hybridization corresponds to a mathematical mixing of
the valence-shell atomic orbitals
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Valence Electron Pair
Geometry
Number of
Orbitals
Hybrid Orbitals
Linear 2 sp
Trigonal Planar 3 sp2
Tetrahedral 4 sp3
Trigonal Bipyramidal 5 sp3d
Octahedral 6 sp3d
2
question : find the hybridazation type and geometry of CH4, PCl5, SF6.
NH3 BeF2
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Combining Atomic Orbitals
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Molecular Orbital Theory
A molecular orbital describes a region of space in a molecule where
electrons are most likely to be found.
Molecular orbitals are formed by combining atomic orbitals on different atoms.
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Bonding in Metals
Electron sea model Band theory.
Bonding in Metals
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Bonding in Metals
Band theory