What the fff

ENGINEERING STUDENT COUNCIL

FACULTY OF ENGINEERING

UNIVERSITY OF SANTO TOMAS

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General Chemistry 1st Term Prelims Reviewer AY1516

Unit 1: FUNDAMENTALS OF CHEMISTRY I. Chemistry and Overview

Chemistry – the study of matter, its composition and properties, the changes that matter undergoes, and the energy associated with these changes. MAIN CHALLENGE OF CHEMISTRY: To understand the connection between the macroscopic world that we experience and the microscopic world of atoms and molecules.

A. Reaction of hydrogen and oxygen 1. Two molecules of hydrogen react with one molecule of oxygen to form two molecules of water

2H2 + O2 -> 2H2O 2. Decomposition of water

2H2O -> 2H2 + O2

3. Hydrogen and Oxygen are chemical elements that exist naturally as diatomic (two-atom) molecule. B. Problem Solving in Chemistry (and life) 1. Making observations (collecting data) 2. Suggesting a possible explanation (formulating a hypothesis) 3. Doing experiments to test the possible explanation (testing the hypothesis) II. The Scientific Method Science is also a plan of action or a procedure for processing and understanding certain types of information. This procedure is called the SCIENTIFIC METHOD. Scientific Method – process that lies at the center of scientific inquiry A. General Framework/Steps in Scientific Method 1. Making Observations Two Types of observation: a. Qualitative – observation based on physical appearance or what is seen by the naked eye; any observation that does not involve a number b. Quantitative – can also be called a measurement, involves both a number and a unit. 2. Formulating hypothesis. Hypothesis is a possible explanation for an observation, or an intelligent guess or solution to the problem observed. 3. Performing experiments. Experiments are done to test a hypothesis. Scientists are able to gather new data that would help validate the hypothesis. Experiments always produce new observations. B. Scientific Models

Theory – often called a model, is a set of tested hypothesis that gives an overall explanation of some natural phenomenon.

Difference between Observation, Theory, and Natural Law

Observation Theory Natural Law

- Something that is witness and can be recorded

- an INTERPRETATION - a possible explanation of why nature behaves in an particular way - possible to change as more information becomes available - educated guess - explanation of behavior; attempt to explain why it happens

- statement formulated from a generally observed behavior - summary of observed (measurable) behavior; a summary of what happens




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C. Units of Measurement Measurement – a quantitative observation that has two parts: a number and a scale (called a unit). "The number without the units is worthless!" 1. Two major systems: a. English system – used in US b. Metric system – used by most countries 2. SI system - International System (le Système International in French); This system is based on the metric system and units derived from the metric system.

The Fundamental SI Units

3. SI Prefixes

The Prefixes used in the SI system

4. Mass vs. weight a. Mass - is a measure of the resistance of an object to a change in its state of motion. b. Weight – the force that gravity exerts on an object to measure its mass weight = (mass)(gravitational constant)




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D. Uncertainty in Measurements 1. Recording Measurements (Significant figures) a. Record all digits that are certain – this are the digits that remains the same regardless o who makes the measurement. b. Record the first digit that is uncertain (all measurements have some degree of uncertainty) c. Uncertainty in the last number (estimated number) is ± 1, unless otherwise noted 2. Accuracy - The agreement of a particular value with the accepted value 3. Precision - The degree of agreement among several measurements made in the same way; refers to the reproducibility of a given type of measurement "You can be precise, but not accurate. If you are accurate, you are necessarily precise."

(a) neither accurate or precise (b) precise but not accurate (c) Both accurate and precise

4. Errors a. Random Errors (indeterminate errors) - Measurements may be high or low; occurs due to 1) Interpretation of the uncertain digit, and 2) Procedural ineptness b. Systematic Errors - Always occur in the same direction, always high or always low; caused by poor measurement calibration E. Significant Figures and Calculations 1. Rules for Counting Significant Figures

Number Rule Example

Nonzero integers Always significant 6.43 m (3 sig. figs.)

Leading zeroes Never significant 0.00643 (3 sig. figs.)

Captive zeroes Always significant 6.0043 (5 sig. figs.)

Trailing zeroes Significant if after a decimal point

63400 (3 sig figs) 0.63400 (5 sig figs)

Scientific notation All digits are significant 6.3400 x 106 (5 sig figs)

2. Multiplication and Division - Keep as many sig figs in your answer as are in the piece of data with the least number of sig figs 2.37 cm x 15.67 cm x 7.4 cm = 274.82046 (keep two sig figs) = 2.7 x 102 cm3 3. Addition and Subtraction - Keep the same number of decimal places as the least precise measurement in your calculation 34.039 m + 0.24 m + 1.332 m + 12.7 m = 48.311 m (keep one decimal place) = 48.3 m 4. Rules for Rounding a. Round at the end of a series of calculations, NOT after each step b. Use only the first number to the right of the last sig fig to decide whether or not to round i. Less than 5, the last significant digit is unchanged ii. 5 or more, the last significant digit is increased by 1




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F. Dimensional Analysis Dimensional Analysis or unit factor method – best way to convert units, you arrive in your answer thru cancellation of units EXAMPLE: Consider a pin measuring 2.85 centimeters in length. What is its length in inches?

1. Unit Conversions Questions a. What units am I given? b. What units must be in my answer? c. What is conversion factor? 2. Full credit can never be given for working a problem in which you do not do all of the following: a. Observe significant figures rules b. Label all steps of your work with the correct units c. Correctly label and identify your answer d. Solve the problem in a manner that can be understood by the reader. 3. Conversion Factors

G. Temperature 1. Three systems of measuring temperature: a. Celsius scale – used in physical sciences b. Kelvin scale – also used in physical sciences c. Fahrenheit scale – used in engineering sciences 2. Unit of temperature = DEGREE 3. Conversion a. Celsius (°C) and Kelvin (K) Kelvin = Celsius + 273.15 Celsius = Kelvin - 273.15 b. Fahrenheit

TC = (TF - 32°F)x (𝟓°𝑪

𝟗°𝑭)

TF = TC x (𝟗°𝑭

𝟓°𝑪)+ 32°F

*Celsius and Fahrenheit is equal at -40o EXAMPLE If a normal body temperature is 98.6oF, what is the normal body temperature in Kelvin? Sol’n: There is no direct conversion from Fahrenheit to Kelvin, you must convert oF to oC first.

TC = (98.g - 32°F)x (5°𝐶

9°𝐹) = 37oC

Kelvin = 37 + 273.15 = 310.15oF




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H. Density Density – the mass(m) of substance per unit volume(v) of the substance

D = 𝒎

𝒗

EXAMPLE Calculate the density of a substance if 643 grams of it occupies 87.3 cm3

D = 𝑚

𝑣 =

643 𝑔

87.3 𝑐𝑚3= 𝟕. 𝟑𝟕

𝒈

𝒄𝒎𝟑

I. Classification of Matter 1. Matter - Anything that occupies space and has mass 2. States of Matter a. Solids - rigid, fixed volume and shape b. Liquids - definite volume, no specific shape c. Gases - no fixed volume or shape, highly compressible 3. Mixtures - Matter of variable composition a. Heterogeneous mixtures - Having visibly distinguishable parts b. Homogeneous mixtures (solutions) - Having visibly indistinguishable parts 4. Components of Mixtures can be Separated by Physical Means a. Distillation b. Filtration c. Chromatography 5. Pure substances a. Elements - Cannot be decomposed into simpler substances by physical or chemical means b. Compounds - Constant composition; Can be broken into simpler substances by chemical means, not by physical means




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J. Fundamental Chemical Laws

LAW Statement Translation Proponent

Law of Conservation of Mass

Mass is neither created nor destroyed. (in a chemical reaction)

In ordinary chemical reactions, the total mass of the reactants is equal to the total mass of the products

Antoine Lavoisier, French chemist; Father of modern chemistry, regarded measurement as the essential operation of chemistry

Law of Definite Proportion

A given compound always contains exactly the same proportion of elements by mass.

Compounds have an unchanging chemical formula

Joseph Proust, French chemist

Law of Multiple Proportions

When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.

Sometimes two elements can come together in more than one way, forming compounds with similar, though not identical formulas

John Dalton; English school teacher

K. Dalton’s Atomic Theory - published by Dalton in 1808, A New System of Chemical Philosophy

Dalton’s Atomic Theory

1. Each element is made up of tiny particles called atoms. 2. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. 3. Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms. 4. Chemical reactions involve reorganization of the atoms—changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

Atomic Masses - atomic weights where mass is often determined by comparison to a standard mass – a process called weighing Avogadro’s Hypothesis - At the same temperature and pressure, equal volumes of different gases contain the same number of particles. L. Early Experiments to Characterize the Atom 1. The Electron and J.J Thomson J.J Thomson – English physicist, studied electrical dischargers in partially evacuated tubes called cathode-ray tubes (the “ray” comes out from the negative electrode or cathode). Thomson proposed that the ray was a stream of negatively charged particles, now called electrons. a. Determined the charge to mass ratio of the electron

𝑒

𝑚= −1.76 × 108 𝐶/𝑔 where e = charge of electron in Coulumbs (C)

m = electron mass in grams




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b. Reasoned that all atoms must contain electrons c. Reasoned that all atoms must contain positive charges d. plum pudding model - model that Thomson proposed in which an atom consisted of a diffuse cloud and positive charge with the negative electrons embedded randomly in it like raisins (electrons) dispersed in a pudding (the positive charge cloud) 2. Robert Millikan and the Oil Drop a. Oil drop experiments determined the charge on an electron b. With charge information, and Thomson's charge/mass ratio, he determined the mass of an electron (9.11x10-31kg) 3. Radioactivity a. in 19th century, scientists discovered that certain elements produce high-energy radiation b. Henri Becquerel – In 1896, he discovered accidentally that a piece of mineral containing uranium could produce its image on a photographic plate in the absence of light and called this phenomenon (the spontaneous emission of radiation by uranium) as radioactivity. c. Three Types of radioactive emission:

1. Gamma (γ) rays - high energy “light” 2. Beta (β) particles - high speed electrons 3. Alpha (α) particles - nuclear particle with a 2+ charge, a charge twice that of the electron and with the opposite sign; mass of an particle is 7300 times that of the electron

4. The Nuclear Atom a. Ernest Rutherford's Metal Foil Experiment 1. Most alpha particles pass straight through thin metal foil because the atom in mostly open space. 2. Some particles were greatly deflected ("like a howitzer shell bouncing off of a piece of paper") i. Could not have been deflected by electrons or single protons ii. Must have been deflected by a positively charged object of substantial mass 3. Disproved Thomson's "plum pudding" model b. Nuclear atom – an atom with a dense center of positive charge (the nucleus) with electrons moving around the nucleus at a distance that is large relative to the nuclear radius.




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M. The Modern View of Atomic Structure: An Introduction 1. Atomic Structure - it consists of a tiny nucleus (with a diameter of about 10-13 cm) and electrons that move about the nucleus at an average distance of about 10-8 cm from it. 2. Nucleus – one of the components of an atom and is small in size (D =10-13 cm) compared to the overall size of an atom; extremely high density; It accounts for almost all the atom’s mass COMPONENTS: a. Protons –POSITIVE CHARGE, equal in magnitude to the electron’s negative charge b. Neutrons – virtually the same mass as a proton but has NO CHARGE 3. Electron – NEGATIVE CHARGE; constitute most of the atomic volume a. The number and the arrangement of electrons causes different atoms to have different chemical properties b. the number of electrons possessed by a given atom greatly affects its ability to interact with other atoms

Particle Mass Charge

Electron 9.109 x 10-31 kg - 1.60 x 10-19 C

Proton 1.673 x 10-27 kg + 1.60 x 10-19 C

Neutron 1.675 x 10-27 kg None

4. Mass and Atomic Number a. Atoms has no net charge, so NUMBER OF PROTONS = NUMBER OF ELECTRONS b. Atomic Number, Z – number of protons c. Mass Number, A – the total number of protons and neutrons d. isotopes - atoms with the same number of protons but different numbers of neutrons; most elements contain mixtures of isotopes

𝐴 𝑍 𝐴

EXAMPLE

Atomic Number Mass Number No. of Proton No. of Electron No. of Neutron

Z A = Z = Z = A - Z

𝑆𝑒3479

34 79 34 34 45

Mass Number

Atomic Number




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Unit 2: ATOMIC STURCTURE AND PERIODICITY The Quantum Mechanical Model of the Atom 1. De Broigle’s Matter-wave – electron bound to an atom behave both as a particle and a wave. 2. Heisenberg Uncertainty Principle - It is impossible for us to know simultaneously the exact momentum and the exact position of a particle in space. The more precisely the position of the particle is measured or determined, the less precisely its momentum can be known, and vice versa. 3. Schrodinger’s Equation

Ĥψ = Eψ Where, ψ = wave function, a function of the coordinates of electron’s position in three-dimensional space. A specific wave function is called an orbital. Ĥ = operator, set of mathematical terms that produce the total energy of the atom when they are applied to the wave function. E = the total energy of the atom (the sum of the potential energy due to the attraction between the proton and electron and the kinetic energy of the moving electron) 4. Quantum (wave) mechanical model a. 1s orbital - wave function corresponding to the lowest energy for the hydrogen atom b. We do not know how the electrons are moving. c. ψ2 = the probability of finding an electron near a particular point in space. The square of the wave function is most conveniently represented as a probability distribution (electron density map; electron density and electron probability), in which the intensity of color is used to indicate the probability value near a given point in space; the darkness of a point indicates the probability of finding an electron at that position

Fig _. The probability distribution for the hydrogen 1s orbital in three-dimensional space

d. radical probability distribution – When the total probability of fi nding the electron in each spherical shell is plotted versus the distance from the nucleus e. For the hydrogen 1s orbital, the maximum radial probability (the distance at which the electron is most likely to be found) occurs at a distance of 5.29 x 10-2 nm or 0.529 Å from the nucleus. i. 1 Å = 10-10 m; the angstrom is most often used as the unit for atomic radius because of its convenient size. Another convenient unit is the picometer: 1 pm = 10-12 m A. Quantum Numbers Quantum Numbers – characterizes the orbitals 1. Principal quantum number (n) = SIZE ANF ENERGY OF THE ORBITAL Represented by: POSITIVE INTERGRAL VALUES – 1, 2, 3,… As n increases, the orbital becomes larger and the electron spends more time farther from the nucleus. An increase in n also means higher energy, because the electron is less tightly bound to the nucleus, and the energy is less negative. 2. Angular momentum quantum number (l) = SHAPE OF ATOMIC ORBITAS, sometimes called as a subshell) Represented by: INTERGRAL VALUES from 0 to n-1 for each value of n

Values of l 0 1 2 3 4

Letters Used s p d f g

3. Magnetic quantum number (ml) = ORIENTATION OF THE ORBITAL IN SPACE RELATIVE TO OTHER ORBITALS Represented by: INTERGRAL VALUES from – l to l, including 0




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EXAMPLE

PRINCIPAL, n

ANGULAR MOMEMTUM, l

CORRESPONDING LETTER of l

SUBLEVEL DESIGNATION MAGNETIC, ml NUMBER OF

ORBITALS

n =1,2,3,… l = 0 to n-1 *combination of n and corresponding letter

ml = – l to l

*just count the number of ml

1 0 s 1s 0 1

2 0 s 2s 0 1

1 p 2p -1, 0, +1 3

3 0 s 3s 0 1

1 p 3p -1, 0, 1 3

2 d 3d -2, -1, 0, 1, 2 5

4 0 s 4s 0 1

1 p 4f -1, 0, 1 3

2 d 4d -2, -1, 0, 1, 2 5

3 f 4f -3.-2, -1, 0,1, 2, 3 7

B. Orbital Shapes and Energies Orbitals - areas of probability for locating electrons Nodal surfaces, or nodes - orbitals contain areas of high probability separated by areas of zero probability. The number of nodes increases as n increases 1. s Orbital – SPHERICAL SHAPE

Fig _. Representation of 1s, 2s, and 3s orbitals

2. p Orbital – two lobes separated by a node at the nucleus; Each orbital lies along an axis (2px, 2py, 2pz); Occur in levels n=2 and greater

Fig _. The electron probability distribution for a 2p orbital.

3. d Orbitals - Occur in levels n=3 and greater Two fundamental shapes: a. Four orbitals with four lobes each, centered in the plane indicated in the orbital label (dxz, dyz, dxy, and dx2-y2) b. Fifth orbital is uniquely shaped - two lobes along the z axis and a belt centered in the xy plane (dz2)




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a) b) Fig _. The boundary surfaces of all five 3d orbitals 4. f Orbitals - Occur in levels n=4 and greater; Highly complex shapes; Not involved in bonding in most compounds

Fig _. Representation of the 4f orbitals in terms of their boundary surfaces

5. Orbital Energies a. All orbitals with the same value of n have the same energy - "degenerate orbitals" b. The lowest energy state is called the "ground state" c. When the atom absorbs energy, electrons may move to higher energy orbitals - "excited state" C. Electron Spin and the Pauli Principle 1. Electron spin quantum number (ms) = MAGNETIC properties of the atom Represented by: + ½ (up) or – ½ (down) only - means that the electron can spin in one of two opposite directions a. concept of electron spin - developed by Samuel Goudsmit and George Uhlenbeck, University of Leyden in the Netherlands; a spinning charge produces a magnetic moment, it seemed reasonable to assume that the electron could have two spin states, thus producing the two oppositely directed magnetic moments 2. Pauli exclusion principle - postulate of Wolfgang Pauli: “In a given atom no two electrons can have the same set of four quantum numbers (n, l, ml, ms). “ a. since only two values of ms are allowed, an orbital can hold only two electrons, and they must have opposite spins.




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D. The History of Periodic Table The periodic table was originally constructed to represent the patterns observed in the chemical properties of the elements. 1. Johann Dobereiner (1780 – 1849) - found several groups of three elements that have similar properties, triads e.g. chlorine, bromine, and iodine 2. John Newlands – English elements, in 1864 suggested that elements should be arranged in octaves, based on the idea that certain properties seemed to repeat for every eighth element in a way similar to the musical scale, which repeats for every eighth tone. 3. Julius Lothar Meyer (1830 – 1895) – German chemist; one of the chemists (other is Mendeleev) which conceived the present form of the periodic table 4. Dmitri Ivanovich Mendeleev (1834 – 1907) – Russian chemist; arranged the elements according to their atomic masses. a. In 1872 when Mendeleev first published his table, the elements gallium, scandium, and germanium (which Mendeleev called ekasilicon) were unknown but there were gaps in his table which was later filled by the elements. 5. The only fundamental difference between this table and that of Mendeleev is that it lists the elements in order by atomic number rather than by atomic mass. E. The Aufbau Principle and the Periodic Table Electron Configuration - describes how electrons are distributed among the various orbitals of an atom

ENERGY LEVEL 1s2 # of e- in sub level

(n, principal quantum number) Sub level (s,p,d,f) Electron Configuration Chart

1. Aufbau Principle – “As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals.” a. orbital diagram – graphic representation of element’s configuration; an arrow represents an electron spinning in a particular direction. 2. Hund’s Rule - The lowest energy confi guration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals.




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3. Valance electrons - the electrons in the outermost principal quantum level of an atom; it is important because they are involved in bonding a. The elements in the same group (vertical column of the periodic table) have the same valence electron configuration. Elements with the same valence electron configuration show similar chemical behavior. b. s-block = Group 1A and 1B; p-block = Group 3A – 8A; d-block = transition metals; f-block = Lanthanide and Actinide series c. Main/Representative elements – Group IA to VIIA = the total number of valence electrons for the atoms in these groups, ns2np3

; member of these groups has the same valence electron configuration.

4. Core electrons – inner electrons of atoms Para mas malinaw:




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EXAMPLE

Element Atomic

No. Last

configuration n l ml ms Period Group

No. of valance

e- Classification

S 16 3p4 3 1 -1 -1/2 3 6A 6 non-metal

Mn 25 3d5 3 2 2 +1/2 4 7B 2 metal

Kr 36 4p6 4 1 1 -1/2 4 8A 0 non-metal

Mg (TRY!)

ans. 12, 3s1, 3, 0, 0, -1/2, 3, 2A, 2, metal

F. Periodic Trends in Atomic Properties 1. Ionization Energy - the energy required to remove an electron from an atom a. Ionization energy increases for successive electrons b. Ionization energy tends to increase across a period i. electrons in the same quantum level do not shield as effectively as electrons in inner levels ii. irregularities at half-filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove b. Ionization energy decreases with increasing atomic number within a group i. electrons farther from the nucleus are easier to remove 2. Electron Affinity - the energy change associated with the addition of an electron; measures the ease of an atom to accept an electron a. Affinity tends to increase across a period b. Affinity tends to decrease as you go down in a period i. electrons farther from the nucleus experience less nuclear attraction ii. Some irregularities due to repulsive forces in the relatively small p orbitals 3. Atomic Radius – distance between atoms in chemical compounds a. determined by measuring half of the distance between radii in a covalently bonded diatomic molecule (covalent atomic radii) or in metal atoms (metallic radii)

b. Radius decreases across a period because of increased effective nuclear charge due to decreased shielding c. Radius increases down a group due to addition of principal quantum levels 4. Electronegativity - a chemical property describing an atom's ability to attract and bind with electrons a. From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. b. From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius. 5. Metallic Character - how readily an atom can lose an electron a. From right to left across a period, metallic character increases because the attraction between valence electron and the nucleus is weaker, enabling an easier loss of electrons. b. Metallic character increases as you move down a group because the atomic size is increasing. When the atomic size increases, the outer shells are farther away.




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SUMMARY:

Trend Left to Right,

Across the period Down to Up,

Up the period

Ionization Energy Increasing Increasing

Electron Affinity Increasing Increasing

Electronegativity Increasing Increasing

Atomic Radius Decreasing Decreasing

Metallic Character Decreasing Decreasing

Increasing: Ionization Energy, Electron Affinity Decreasing: Atomic Radius, Metallic Character

EXAMPLE 1. Which is more electronegative, Fluorine or Sodium? ANS. Fluorine 2. Arrange the ff. elements by increasing atomic radius, Antimony, Nickel, Barium, Palladium, and Xenon? ANS. Xenon, Antimony, Nickel, Palladium, Barium 3. Which has the lowest ionization energy among Helium, Yttrium, and Cobalt? ANS. Yttrium

Unit 3: GENERAL CONCEPTS OF CHEMICAL BONDING A. Types of Chemical Bonds Chemical bond - forces that hold groups of atoms together and make them function as a unit; A bond will form if the energy of the aggregate is lower than that of the separated atoms. Bond energy – the energy required to break the bond 1. Ionic Bonding (electron-transfer)- Ionic substances are formed when an atom that loses electrons relatively easily reacts with an atom that has a high affinity for electrons. a. Electrons are TRANSFERRED b. Happens when metals react with nonmetals which results to an ionic compound 2. Coulomb’s Law – calculates the energy of interaction between a pair of ions

𝐸 = (2.31 × 10−19 𝐽 ∙ 𝑛𝑚)(𝑄1𝑄2

𝑟)

where E = energy, J r = distance between the ion centers, nm Q1, Q2 = numerical ion charges. a. If the answer is negative (-), it indicates an attractive force; The ion pair has lower energy than the separated ions. b. can be used to calculate the repulsive energy when two like charged ions are brought together; this will a POSITIVE sign. 3. Bond length - the distance at which the system has minimum energy. a. For proton-electron, Attractive forces work

Increasing: Ionization Energy, Electron Affinity Decreasing: Atomic Radius, Metallic Character




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b. For electron-electron/proton-proton, Repulsive force work c. Energy is given off (bond energy) when two atoms achieve greater stability together than apart 4. Covalent Bonding (electron-sharing) - electrons are shared by nuclei; two identical atoms share electrons equally. The bonding results from the mutual attraction of the two nuclei for the shared electrons. a. polar covalent bonding - the atoms are not so different that electrons are completely transferred but are different enough that unequal sharing result b. non-polar covalent bonding - Electrons are shared evenly c. bond polarity - development of the partial positive (σ+) and negative (σ-) charges on the atoms B. Electronegativity Electronegativity - the ability of an atom in a molecule to attract shared electrons to itself. 1. Linus Pauling – American scientist, developed the most widely accepted method for determining values of electronegativity The difference (Δ) between the actual (measured) and expected bond energies:

Δ = (H – X)act – (H – X)exp

a. If H and X are identical atoms (having same electronegativity), the difference (Δ) is ZERO. The electrons in the bond are shared equally, and no polarity develops. b. if X has a greater electronegativity than H, the shared electron(s) will tend to be closer to the X atom c. The greater is the difference in the electronegativity of the atoms; the greater is the ionic component of the bond and the greater is the value of Δ. Thus the relative electronegativity of H and X can be assigned from the Δ values.

Identical Atoms ZERO electronegativity difference Electrons shared equally

Different atoms Noticeable electronegativity difference Electron transfer occur

C. Bond Polarity and Dipole Moments 1. Dipole Molecules – molecules that said to have a dipole moment a. Dipole moment - when a molecule has a center of positive charge and a center of negative charge - has a partial positive end and partial negative end - any diatomic (two-atom) molecule that has a polar bond and polyatomic molecules can exhibit dipolar behavior b. Representation of charge distribution i. Arrow and tail – the arrow pointing to the negative charge center and the tail indicating the positive charge center




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ii. Electrostatic potential diagram – Variation in charge distribution is represented by colors of visible light. Red indicates the most electron-rich region (negative) of the molecule and blue indicates the most electron-poor region (positive).

EXAMPLE

No Dipole Moment – CO2 With Dipole Moment – H2O

TYPES OF MOLECULES WITH POLAR BONDS BUT NO RESULTIONG DIPOLE MOMENT

D. Ions: Electron Configuration and Size 1. Bonding and Noble Gas Electron Configurations a. Ionic bonds - Electrons are transferred until each species attains a noble gas electron configuration b. Covalent bonds - Electrons are shared in order to complete the valence configurations of both atoms 2. Predicting Formulas of Ionic Compounds Ionic compounds – usually referring to the solid state of that compound since at solid state, ions are close together a. To predict the formula of the ionic compound, we simply recognize that chemical compounds are always electrically neutral—they have the same quantities of positive and negative charges.

CaO




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Common Ions with Noble Gas Configuration in Ionic Compounds

Group 1A Group 2A Group 3A Group 6A Group 7A Electron Configuration (of Noble gas the ions would like to achieve = lose 1 e- = lose 2 e- = lose 3 e- = gain 2 e- = gain 1 e-

H-, Li+ Be2+ [He]

Na+ Mg2+ Al3+ O2- F- [Ne]

K+ Ca2+ S2- Cl- [Ar]

Rb+ Sr2+ Se2- Br- [Kr]

Cs+ Ba2+ Te2- I- [Xe]

b. Take note of this rule: Ions generally adopt noble gas electron configurations in ionic compounds. c. Formulas for compounds are balanced so that the total positive ionic charge is equal to the total negative ionic charge

Ex. 𝐴𝑙2+3𝑂3

−2 Total positive = +6; Total Negative = -6 3. Sizes of Ions a. Anions are larger than the parent atom b. Cations are smaller than the parent atom c. Ion size increases within a family d. Isoelectronic ions i. Ions with the same number of electrons ii. Size decreases as the nuclear charge Z increases E. Lewis Structures 1. Lewis Structure of molecules – show how the valence electrons are arranged among the atoms in the molecule 2. Most important requirement for the formation of a stable compound is that the atoms achieve noble gas electron configurations. 3. In writing Lewis structures, only the valence electrons are included. Use dots to represent the electrons 4. Duet rule - Hydrogen, lithium, beryllium, and boron form stable molecules when they share two electrons (helium configuration) 5. Octet rule - Elements (carbon and beyond) form stable molecules when they are surrounded by eight electrons (noble gas configuration) EX. F2 molecules

Determine the number of a) Bonding pair = 1 (pertains to the molecules shared electrons) b) Lone pair/s = 3 (pertain to electrons not involved in bonding)

5. Writing Lewis Structures a. Add up the TOTAL number of valence electrons from all atoms b. Determine the central atom.

*Tip: Kapag may carbon ung compound na pinapadrawing, kadalasan un ung central atom. Or minsa ung element na may pinakamadaming bonds ka na maiikakabit

c. Use a pair of electrons to form a bond between each pair of bound atoms. Lines are used to indicate each pair of bonding electrons. Dots are used to represent lone pairs

*mas maganda siguro kung gawa ka muna ng single bonds, tpos kapag may mga kulang tsk aka magadd ng double bond

d. Arrange the remaining atoms to satisfy the duet rule for hydrogen and the octet rule for the second row elements




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*always double check your structure, count the number of electron of your structure (for each bond, 2e- and for each dot, 1e-) and make it sure it is equal to the calculated total number of valence

EXAMPLE Draw the Lewis structure of HNO3 (a) Determine the number of valence electron: H: 1 x 1e- = 1 N: 1 x 5e- = 5 O: 3 x 6e- = 18 24e-

(b) Determine number of electron to achieve octet/duet rule: H: 1 x2- = 2 N: 1 x 8e- = 8 O: 3 x 8e- = 24 34e-

(c) Number of bonding electrons: No. of octet e-- - No of valance = 34 – 24 = 10e-

(d) Number of bonds: No. of bonding e-/2 = 10/2 = 5 bonds (e) Nonbonding e-: No. of valence e- - No. of bonding e-- = 24 – 10 = 14e-

you can now construct the structure: O = N – O – H O – N – O – H

O (resonance could O

occur thru e- delocalization)

F. Exceptions to the Octet Rule 1. Boron Trifluoride a. Note that boron only has six electrons around it; total of 24e= b. BF3 is electron deficient and acts as a Lewis acid (electron pair acceptor) c. Boron often forms molecules that obey the octet rule 2. Sulfur Hexafluoride a. Note that sulfur has 12 electrons around it, exeeding the octet rule b. Sulfur hexafluoride is very stable c. SF6 fills the 3s and 3p orbitals with 8 of the valence electrons, and places the other 4 in the higher energy 3d orbital 3. More About the Octet Rule a. Second row elements C, N, O and F should always obey the octet rule b. B and Be (second row) often have fewer than eight electrons around them, and form electron deficient, highly reactive molecules c. Second row elements never exceed the octet rule, since their valence orbitals (2s and 2p) can accommodate only eight electrons. d. Third row and heavier elements often satisfy (or exceed) the octet rule e. Satisfy the octet rule first. If extra electrons remain, place them on elements having available d orbitals i. When necessary to exceed the octet rule for one of several third row elements, assume that the extra electrons be placed on the central atom G. Resonance Resonance - When more than one valid Lewis structure can be written for a particular molecule. The resulting electron structure of the molecule is given by the average of these resonance structures. 1. Nitrate ion - Experiments show that all N-O bonds are equal a. A single Lewis structure cannot represent the nitrate ion




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b. A resonance structure is drawn by writing the three variant structures, connected by a double-headed arrow (the arrow shows show that the actual structure is an average of the three resonance structures” c, The correct description of NO3

` is not given by any one of the three Lewis structures but is given only by the superposition of all three. 2. Odd Electron Molecules - Molecules in which there is not an even number of electrons; does not fit localized electron model 3. Formal Charge - Number of valence electrons on the free atom minus Number of valence electrons assigned to the atom in the molecule FC = no. of valance e- - no. of lone pairs – ½ (no. of shared or bonded e-) a. Lone pair (unshared) electrons belong completely to the atom in question b. Shared electrons are divided equally between the sharing atoms c. The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species d. If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the bonding in the molecule or ion. H. Molecular Structure: The VSEPR Model Molecular structure - the threedimensional arrangement of the atoms in a molecule. 1. Valence Shell Electron Pair Repulsion (VSEPR)-. The structure around a given atom is determined principally by minimizing electron-pair repulsions;. Non-bonding and bonding electron pairs will be as far apart as possible a. Lone (unshared) electron pairs require more room than bonding pairs (they have greater repulsive forces) and tend to compress the angles between bonding pairs b. Lone pairs do not cause distortion when bond angles are 120° or greater 2. VSEPR and Multiple Bonds a. For the VSEPR model, multiple bonds count as one effective electron pair b. When a molecule exhibits resonance, ANY of the resonance structures can be used to predict the molecular structure using the VSEPR model 3. Molecules Containing No Single Central Atom a. Apply the principal of distancing shared and unshared electron pairs 3. Look at real 3-dimensional, rotatable models to develop predictive skills




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Arrangement of Electron Pairs Around an Atom Yielding Minimum Repulsion A = Central Atom; X = bonded pairs, E = lone pairs

No. of electron

pairs

Bond pairs

Lone pairs

FORM NAME STRUCTURE ANGLES

2 2 0 AX2 Linear

1800

3 3 0 AX3 Trigonal Planar

120o

3 2 1 AX2E Bent (V shaped)

<120o

4 4 0 AX4 Tetrahedral

109.50

4 3 1 AX3E Trigonal pyramidal

<109.5o

4 2 2 AX2E2 Bent (V shaped)

<109.50

5 5 0 AX5 Trigonal bipyramidal

90o 120o

5 4 1 AX4E Seesaw

<90o <120o

5 3 2 AX3E2 T-shaped

<90o




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5 2 3 AX2E3 Linear

180o

6 6 0 AX6 Octahedral

90o

5 5 1 AX5E Square pyramidial

<90o

4 4 2 AX4E2

Square Planar

<90o




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Unit 4: STOICHIOMETRY OF CHEMICAL FORMULAS AND EQUATIONS A. Counting by weighing 1. Get the average mass of objects. a. For purposes of counting, the objects behave as though they were all identical, as though they each actually had the average mass. b. to relate the mass to a number of atoms, we must know the average mass of the atoms. B. Atomic Masses 1. C12, the Relative Standard – C12 is assigned a mass of exactly 12 atomic mass units (amu) a. Masses of all elements are determined in comparison to the carbon - 12 atom (12C) the most common isotope of carbon b. Comparisons are made using a mass spectrometer (most accurate method currently available for comparing the masses of atoms) 2. Atomic Mass (Average atomic mass, atomic weight) a. Atomic masses are the average of the naturally occurring isotopes of an element b. Atomic mass does not represent the mass of any actual atom c. Atomic mass can be used to "weigh out" large numbers of atoms C. The Mole Mole - the number equal to the number of carbon atoms in exactly 12 grams of pure 12C 1. Avogadro's number - One mole of something consists of 6.022 x 1023 units of that substance 6.022 x 1023 units = 1 mole 2. Measuring moles 1. An element's atomic mass expressed in grams contains 1 mole of atoms of that element 2. 12.01 grams of carbon is 1 mole of carbon 3. 12 grams of carbon-12 is 1 mole of carbon-12 EXAMPLE A silicon chip used in an integrated circuit of a microcomputer has a mass of 5.68 mg. How many silicon (Si) atoms are present in the chip?

Ans. 5.68 𝑚𝑔 𝑆𝑖 × 1 𝑔 𝑆𝑖

1000 𝑚𝑔 𝑆𝑖 ×

1 𝑚𝑜𝑙 𝑆𝑖

28.09 𝑔 𝑆𝑖 ×

6.022 × 1023 𝑎𝑡𝑜𝑚𝑠

1𝑚𝑜𝑙 𝑆𝑖= 𝟏. 𝟐𝟐 × 𝟏𝟎𝟐𝟎 𝒂𝒕𝒐𝒎𝒔

*gamitin ang factor-label method. Gawing clues ang mga units ng given mo, dapat magcancel ung same units hanggang matira ung desired mo. D. Molar mass Molar Mass (Gram molecular weight) - The mass in grams of one mole of a compound. The sum of the masses of the component atoms in a compound Ex. Molar mass of ethane (C2H6): Mass of 2 moles of C = 2(12.01 g) Mass of 6 moles of H = 6(1.008 g) 30.07 g




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E. Percent Composition of Compounds 1. Calculating any percentage - "The part, divided by the whole, multiplied by 100" 2. Percentage Composition - Calculate the percent of each element in the total mass of the compound Mass Percent/Weight Percent = (#atoms of the element)(atomic mass of element) x 100 (molar mass of the compound) EXAMPLE Mass percent of each element of Ethanol, C2H5OH

Element Number of

Moles Molar Mass,

g/mol Mass (Mol x MW)

Mass Percent

(𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡

𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 × 100)

C 2 12.01 2 mol C x 12.01 𝑔

𝑚𝑜𝑙= 24.02 52.14%

H 6 1.008 6.048 13.13%

O 1 16.00 16.00 34.73%

SUM = 46.07 g 100%

F. Determining the Formula of a Compound

1. Determining the empirical formula a. Determine the percentage of each element in your compound b. Treat % as grams, and convert grams of each element to moles of each element c. Find the smallest whole number ratio of atoms d. If the ratio is not all whole numbers, multiply each by an integer so that all elements are in whole number ratio 2. Determining the molecular formula a. Find the empirical formula mass b. Divide the known molecular mass by the empirical formula mass, deriving a whole number, n c. Multiply the empirical formula by n to derive the molecular formula EXAMPLE Determine the empirical and molecular formulas for a compound that gives the following percentages on analysis (in mass percent): 71.65% Cl; 24.27% C; 4.07% H The molar mass is known to be 98.96 g/mol.

Soln:

BY Empirical

Basis = 100 g

Element Percentage Mass (%*100) Moles

Cl 71.65% 71.65 71.65 g Cl x 1 𝑚𝑜𝑙 𝐶𝑙

35.45 𝑔 𝐶𝐿= 2.021 1

C 24.27% 24.27 2.021 1

H 4.07% 4.07 4.04 2

Therefore, CClH2

BY molecular

Empirical formula mass = 49.48 g/mol Molar mass is given = 98.96 g/mol

Molar mass = 98.96 g/mol = 2 Empirical formula mass 49.48 g/mol

Molecular formula= (ClCH2)2 = C2Cl2H4 = dichloroethane.

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