CHEMISTRY FOR ENGINEERS Flashcards
PRELIMINARY EXAMS 1ST YEAR (93 cards)
1
Q
Aggressive of at least two atoms in a definite arrangement held by chemical force
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Atom
2
Q
CLASSIFICATION OF MATTER
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MIXTURE PURE SUBSTANCE
3
Q
Combination of two or more substances in which cannot be
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Mixture
4
Q
Two types of Mixture
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Homogeneous Heterogeneous
5
Q
Stirring
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Homogenous
6
Q
Solution mixture of two or more substances
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Homogeneous
7
Q
Suspension
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Heterogeneous
8
Q
dispersed into another particles
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Heterogeneous
9
Q
Has a definition or constant composition of distinct properties
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Pure substance
10
Q
Two types of Pure substance
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ELEMENTS COMPOUND
11
Q
Two elements chemically bonded
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Compound
12
Q
Is an atom of substance
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Element
13
Q
Can be separated by chemical method
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Element and Compound
14
Q
Enough seen not settle down
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Colloid
15
Q
Very high temperature of stars
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Plasma
16
Q
Physical quantity Unit and symbol
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Time, second, s
Temperature, kelvin, K
Degree, C
Fahrenheit Fdegree
Lenght, meter, m
Argstorm, A
Mass, Atomic mass unit, atomic mass unit/u
Pound, lb
Kilogram, Kg
Amount of substance, mol, mol
Electric current, ampere, A
Illuminous Intensity, Candela, Cd
Area, Square meter, m^2
Hectare, ha
Square yard, yd^2
Volume, Liter, L
Cubic centimeter, cm^3
US galoon, gal
Density, Kg per cubic meter, Kg/m^3
Gram per cubic, g/cm^3
17
Q
Absolute deviation
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Gives the exact amount by which measurement deviates from the mean
18
Q
Formula for absolute deviation
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Absolute deviation . mean of measurement/100
19
Q
Precision
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Refers to how close measurement of the same item to are to each other
20
Q
Accuracy
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How refers to how close a measurement is to the accepted or true value
21
Q
Absolute error
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This is the differences between the measure value and true value
22
Q
Formula for absolute error
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Absolute error = measured value - true value
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Q
Formula for relative error
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Relative error = Measured value/absolute error - True value/ True value often x to 100
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Q
Measure of the accuracy of measurement calculation, It expresses the magnitude of the error ( the difference between the measured value and the true value) relative to the formula
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Relative error
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Fields of Chemistry
Analytical chemistry
Biochemistry
Inorganic chemistry
Organic chemistry
Polymer chemistry
Medicinal chemistry
Physical chemistry
Theoretical chemistry
Environmental chemistry
Industrial chemistry
Materials chemistry
Nuclear chemistry
Geochemistry
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The study of the composition of substances. This field involves identifying and quantifying materials in a mixture, determining the structure of compounds and analyzing chemical properties
Analytical chemistry
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The study of chemical processes within and related to living organism. This field combines biology and chemistry to understand cellular process at molecular level
Biochemistry
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the study of inorganic compound, which include minerals, metals, and non metals, excluding most organic ( carbon based) compounds.
Inorganic compounds
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The study of carbon- containing compounds, including their structure, properties, reactions, and synthesis.
Organic chemistry
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The study of how matter behaves on molecular and atomic level and how chemical reactions occur, this field combines principle of physics and chemistry
Physical chemistry
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The use of mathematics and computer simulations to understand chemical systems and predict the outcomes of chemical reactions
Theoretical Chemistry
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The study of chemical process occurring in the environment, including the effects of human activity on the environment and how to remediate pollution
Environmental chemistry
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The application of chemical process to the manufacture of products on a commercial scale, such as chemicals, materials, and pharmaceuticals.
Industrial chemistry
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The study of the chemistry of polymers, large molecules made up of repeating units, this includes understanding their synthesis, structure, properties, and applications.
Polymer chemistry
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The study of the design and properties of materials, particularly solids, It includes the study of metals, ceramics, polymers, and composites.
Material chemistry
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The study of chemistry as it relates to the design, synthesis, and development of pharmaceutical agents (drug)
Medicinal chemistry
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The study of radioactive elements and their reactions, Including nuclear reactions, radioactive decay, and application in energy production and medicine
Nuclear chemistry
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The study of the chemical composition of the Earth and other planets, Including the distribution and movement of chemical elements and compounds
Geochemistry
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Are the digits in a number that carry meaningful information about precision. They indicate the accuracy of a measurement or calculation.
Significant figures
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Significant figures
Non- zero digits; all non zero digits (1-9) are always significant example 123 has significant figures
Zeros between non zero digits are significant example 1002 has four significant figures
Leading zeros, Zeros to the left of the first non zero are not significant, they are only place holders. Example 0.0025 has two significant figures
Trailing zeros with a decimal point, zeros to the right of a non zero digit in a decimal number are significant. Example 2.300 has four significant numbers.
Trailing zeros without a decimal point. zeros to the right of a non zero digin in a whole number without a decimal point may or may not be significant. it depends on whether theres a way to indicate that the zeros are measured
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The Total mass of substance present after a chemical reaction is the same as the total mass of substance before the reaction.
Law of conservation of mass
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Total mass of substances after a chemical reaction = total mass of substance before the reaction
Law of conservation of mass
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A given compound always contains exactly the same proportion of elements by mass
Law of definite proportion
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All samples of a compound have the same composition the same proportions by mass of the constituent elements.
Joseph proust
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Law of multiple proportions
John dalton
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Is an english teacher and a greek atomist. He is the propornents of daltons atomic theory
John dalton
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Each chemical element is composed of minute, indivisible, particles called atoms. Atoms can neither be created or destroyed during a chemical change
All atoms of the same elements are alike in mass (weight) and other properties but the atoms of one element are different from those all of other elements
Atoms are indivisible. We cannot add atoms in fractions
Atoms are indestructible after chemical masses are unchanged. ( Consistent with the Law of Conservation of mass
Dalton Atomic theory
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When two elements form a series of compounds, the ratios of the masses of the second elements that combine with 1 gram of the first element can always be reduced to small whole numbers
LAW OF MULTIPLE PROPORTIONS
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Each element is made up of tiny particles called atom.
The atoms of the given element are identical, The atoms of the different elements are different in some fundamental ways or ways.
Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relatives numbers and types of atoms.
chemical reactions involve reorganization of the atoms themselves are not changed in the chemical reaction
Dalton atomic theory
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EARLY MODELS AND EXPERIMENTS TO CHARACTERIZED THE ATOM
DALTON SOLID SPHERE MODEL
THOMPSONS PLUM PUDDING MODEL
RUTHERFORDS NUCLEAR MODEL
BOHRS SOLAR SYSTEM MODEL
QUANTUM MECHANICAL MODEL
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Was the first atomic model and was developed by john dalton in 1808. He hypothesized that an atom is a solid sphere that could not be divided into smaller particle.
Held the theory of structureless atom (Billiard ball model)
Dalton's Solid sphere Model
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Hypothesized that the hydrogen atom is fundamental.
All other elements made up of hydrogen atoms
His hypothesis was rejected by the 1830s for (example chlorine atom had mass of 34.5 times of that hydrogen)
William prout
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rasin bread model or plum pudding model of the atom the electrons were embedded in a uniform sphere of positive charge like blueberries stuck into a muffin.
All atoms are made up of a combination of positive and negative particles.
Thompson's plum pudding model
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Who discovered that the charge of electron is e = -1.60x10^-19c and the charge to mass ratio is equal to -1.76x10^8
ROBERT MILLIKAN
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The mass of the atom is concentrated in a small volume at the center called nucleus.
The nucleus accounts for more than 99% of the mass of the atom.
Diameter of an atom is 1x10^-10cm
Diameter of nucleus is approximately 1x10^13cm
Nucleus is positively charged
Rutherford's Nuclear model
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The model is also referred to as the planetary model of an atom.
If electrons orbit the nucleus in set orbits, each with quantum number n , this can be used to explain the hydrogen atomic spectrum.
Since electrons energy is quantized they can only absorb or emit photons with energy equal to the difference in permissible orbits.
Bohrs Solar system Model
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Dual nature of light
Electromagnetic radiation (light) has particulate nature light can be viewed as a stream of photons. Each photons carries a packet of energy called a quantum and this energy can be calculated using this equations. And what are those equations?
E = hv where E = energy
h = is Planck's constant 6.626x10^-36 J-s
v = Frequency of light (Hz)
or by this equation
E = hc/ wavelength of light where c = speed of light
^ = wavelength of light
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Erwin schroeidinger
Wave model
Based in the duality of light and matter.
Electrons can be thought of as waves and both their energy and behaviour can be thought of as wave function that depend on parameters known as the quantum numbers
Quantum Mechanical Model's
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EVIDENCES OF ATOMS
Robert Boyle's experiment on bases which led to Boyle's law (1661).
Antoine lavoisier's experiment that led to the law of conservation of mass.
Joseph louis proust experiment that led to the Law of constant proportion ( Law of definite proportion).
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FUNDAMENTAL PARTICLES OF AN ATOM
Protons: Positively charge, in nucleus
Neutron: No charge, In nucleus
Electrons: Negatively charge, orbit nucleus
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Mass of proton = Neutron > electron
Charge: proton = +1, Neutron = 0 electron = -1
Fundamental of an atom
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Key atomic concepts
Atom: smallest unit of matter that retains properties of an element
Atomic number : Number of protons in an atom's nucleus
Mass number: total number of protons and neutrons in an atom's nucleus.
Isotopes: atoms of the same element with different mass, the same no. of protons and different no. of neutrons.
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Parts of atom
ELECTRON PROTON NEUTRON
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Discovered by JJ thompson: name given by george stoney in 1874
Relative mass (amu) = 0.00054868
Mass in grans = 9.11x10^-28
Actual charge = -1.6x10^-19 C
Assign charged = -1
Electron
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Discovered by rutherford: name given by Goldstein
Relative mass (amu) = 1.007280
Mass in grans = 1.67x10^-24
Actual charge = +1.67x10 ^-19 C
assigned charge = +1
Proton
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Discovered and named by james chadwick, 1932
Relative mass (amu) = 1.008670
Mass in grans = 1.67x10^-24
Actual charge = 0
assigned charge = 0
Neutron
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Energy levels : Main shells around the nucleus (n=1,2,3....)
Sublevels: Subdivisions of energy levels (s ,p , d, f)
Orbitals: specific regions where electrons are likely to be found
Relationship:> Energy levels > sublevels > orbitals
Energy levels Subslevel orbitals relationship
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Is a region in space of some general distance from the nucleus where a group of electrons is most likely to be found
Energy level
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Is a subdivision of an energy level in an atom: made up of set of orbital
Sub level
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Sublevel and their max electrons
Sublevel Max number of electrons
S = 2
p = 6
d = 10
f = 14
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Is the region in space, within an atom or molecule where there can be no more than two electrons
Sub Level Maximum number of electrons Maximum orbitals
S = 1 , 1
p = 6 -, 3
d = 10, 5
f = 14, 7
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States that the orbitals are filled in order of increasing energy
Aufbau Principle
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Specifies that orbital should be filled singly before pairing
Hund's rule
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Specifies that the maximum number of electrons which an orbital may accommodate is two
Pauli Exclusion Principle
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Based on quantum theory and wave mechanics
Describes electrons as waves rather than particles
Introduces concept of probability of finding an electron
Uses quantum numbers to describe electron states
Quantum Mechanical model
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Describes the main energy level or shell that an electron occupies
Principle quantum number = n
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Designates the shape of region in space an electron occupies Designates sublevel, specific atomic orbital that an electron may occupy
Angular momentum quantum number
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Designates the specific orbital within a subshell
Magnetic quantum number, ml
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Refers to the spin of an electron orientation of the magnetic field produced by this spin
Spin quantum number, ms
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Drawing lewis structure method
Count the valence of electrons
Draw the skeleton structure
Distribute remaining electrons
Check the octet rule
Add double or tripple bonds if necesseray
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Are the attractive forces that hold atoms together to form molecules and compounds.
The three primary types are ionic, covalent, and metallic bond
Covalent bonds
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Is a measure of how strongly an atom attracts electrons when its in a chemical bond
electronegativity
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Atom share electrons to achieve a full outer electron shell
Tends to be strong and have specific bond length based on the atom involved
Covalent bonds
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Bonds form between atoms with different electronegativity
Polar covalent bond
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occurs when two atoms share electrons equally or nearly equally
Non polar covalent bond
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Refers to the distance from the nucleus of an atom to the outmost electron shell. its measure of how large an atom is
Atomic size
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Is a measure of the size of an ion which is an atom that has gained or lost one or more electrons
Ionic radius
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Losing one or more electrons
Positive ion
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Gaining one or more electrons
Negative ion Anion
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The amount of energy required to remove an electron from an atom or ion in its gaseous state.
Ionization energy
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A measure of how much energy is released when an atom in the gaseous state gains an extra electron to form negative ion
Electron affinity
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Are atoms ions or molecules that have the same number of electrons and therefore the same electron configuration but they may have different numbers protons and different chemical properties
Isoelectronic species
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