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UF / Chemistry / CHM 1025 / Which table houses groups or families of elemeents that sahre comparab

Which table houses groups or families of elemeents that sahre comparab

Which table houses groups or families of elemeents that sahre comparab

Description

School: University of Florida
Department: Chemistry
Course: Introduction to Chemistry
Professor: Melanie veige
Term: Fall 2015
Tags: ions, Molecular, naming, study, guide, chm, and Chemistry
Cost: 50
Name: CHM 1025 Midterm Study Guide Ch 1-4
Description: This study guide covers what is going to be on Exam 1
Uploaded: 09/30/2017
11 Pages 166 Views 2 Unlocks
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Chapter 1 


Which table houses groups or families of elemeents that sahre comparable properties?



Matter – Anything that occupies space and has mass.

Mass – a measure of the quantity of matter.

The interaction of mass with gravity creates weight.

Energy – the capacity to move an object or transfer heat.

Pure Substance – always has the same chemical composition no matter its origin. Mixture – Two or more substances that may vary in composition. - You can separate a mixture into their component pure substances. This  can be done physically, by grinding/dissolving/filtering.

- Homogeneous Mixture AKA Solution – uniform composition throughout.  Most mixtures are dissolved in water and are often clear. Each sample, no  matter the size, would have the same components in the same  proportions.


What refers to a solution in which a substance is dissolved in water?



If you want to learn more check out How do traits influence behavior?

- Heterogeneous Mixture – Not uniform throughout. Different samples  will have different proportions of components.

- If the material can be separated, them it is a mixture, if not, it is a pure  substance.  

- If the substance can be broken down, it is a compound, if not, it is an  element

- If the element conducts electricity, it is a metal, if not, it is a nonmetal. - Not all solutions are liquids. Solid solutions are called alloys.

There are two types of pure substances: Elements and Compounds Element – a substance that cannot be broken down into other simpler substances  even by chemical reaction.

- An atom is the smallest unit of an element that has the chemical  properties of that element.


What is the equation for specific heat?



If you want to learn more check out What is the definition of redlining?

- Molecule – Two or more atoms bound together in a discrete arrangement. Molecules can also be formed by the combination of only one element. Ex) O2 Don't forget about the age old question of What does survival of the fittest state about every state?

- Element symbols with subscripts represent a ratio of elements in a  compound.

Periodic Table – Columns of the table house groups/families of elements that share comparable properties.

- 2 main categories: Metals and Nonmetals  

- Metals – shiny/conductive/ductile/malleable

- Nonmetals – dull/poor conductors/brittle

Compound – a substance made up of two or more elements combined in definite  proportions.

- Compounds have different properties than their component elements. Chemical Formula – describes the composition of a compound using symbols for  the elements that make the compound.

- Subscripts how the relative proportions of the elements in the  compound. If there is no subscript, the subscript is assumed to be one. Physical State – a form that matter can take. Solid/Liquid/Gas Don't forget about the age old question of What are the sub-areas of linguistics?

- Solid – Fixed shape

- has its own volume

- no volume change under pressure

- particles are fixed in place and are in a regular array  

(Crystalline).

- Liquid – takes the shape of its container

- has its own volume

- slight volume change under pressure

- particles are randomly arranged and move/bump into each  

other

- Gas - takes the shape of its container and fills it

- has the volume of the container

- large volume change under pressure

- particles are widely separated and move independently

Symbols for Physical States

- Aqueous (aq)

- Solid (s)

- Liquid (l)

- Gas (g)

Aqueous Solution (aq) – a solution in which a substance is dissolved in water. Solids – Crystalline solids have a regular ordered structure. Amorphous solids have no ordered structure.

Physical Properties – can be measured without changing the identity of the substance and are often easily reversed. Ex) Color/odor/density/temperature/mass/volume/conductivity/malleability/ductility  Chemical Properties – how something reacts with something else and changes the identity of the substance. Ex) flammability/color change/corrosiveness/combustion/reactivity Don't forget about the age old question of What is the state of membrane in the fluid mosaic model?

Intensive Properties – independent of the amount of substance. Ex) Color Extensive Properties – depend on the amount of substance. Ex) Energy released upon combustion/mass/volume/length

Changes of state are always physical.

Energy – Energy Units

- SI unit of hear = Joule (J) – 1J = (1kg X 1m^2)/S^2)

- Kinetic Energy = (1/2)mv^2

- Calorie (cal) – the quantity of energy (heat) needed to raise3 the temperature of 1g of water by 1 degree Celsius

- 1 cal = 4.184J

- 1 kcal = 1000 cal = 1 Cal

Heat and Specific Heat

Heat (q) = the energy transferred between a system and its surroundings as a result of temperature difference.

Specific Heat (s or SH) = the quantity of heat required to change the temperature of a 1g of a substance by 1 degree Celsius or 1-degree Kelvin. Heat (q) = m (mass) X SH X change in temperature (Temp 1 - Temp 2) Units = g X J/(g X degrees C or K) X Degrees C or K Heat Transfer

- m (substance 1) X SH (substance 1) X change in temperature (substance 1) = -m (substance 2) X SH (substance 2) X change in temperature (substance 2)

- Pay attention to the negative sign on the right side of the equation. - If the negative sign is present at the end, this means that heat is released, not absorbed.

Chapter 2 We also discuss several other topics like What are the two features of anarchy?

2.1 – Dalton’s Atomic Theory

- Antoine Lavoisier’s – Law of Conservation of Mass – No measurable  change in mass occurs during chemical reaction - Law of Definite Proportions – samples of the same compound always  contain the same proportions by mass of the component elements. - Dalton’s Atomic Theory

1) All matter is composed of small indivisible particles called atoms. 2) All atoms of a given element are identical in mass and chemical properties, but different elements have different masses and different  chemical properties.

3) Atoms are not created or destroyed in chemical reactions.

4) atoms combine in simple whole number ratios to form compounds. - A chemical reaction is a rearrangement of atoms into a new combination,  resulting in one or more new chemical substances.  

- Qualifications to Dalton’s Theory – The 1st statement is not strictly true, atoms are  composed of smaller particles, subatomic particles, and the 2nd statement  leaves out that atoms of the same element can vary in mass.

2.2 – Structure of the Atom

- Subatomic Particle – Proton, Neutron, Electron

- Electron – Negative Charge of -1.6022 X 10 ^-19 coulombs  

- Proton – Positive Charge of +1.6022 X 10 ^-19 coulombs

- Neutron – Uncharged – accounts for extra mass in the nucleus - To be electrically neutral, atoms must have equal numbers of protons and  electrons.

- The charge of an electron is expressed as a 1- superscript  

- The charge of a proton is expressed as a 1+ superscript

- Hans Geiger and Ernest Rutherford discovered the nucleus through a gold foil  experiment.

- The nucleus contains most of the mass of the atoms and electrons exist outside  the nucleus in an electron “cloud”.

- Atomic Number – the number of protons in an atoms nucleus that determine the  identity of that element. (Z)  

- The atomic number is just above the element symbol on the periodic table. - Not all atoms of the same elements contain the same number of neutrons  - Isotope - an atom of an element that has a specific number of neutrons.

- As long as the number of protons does not change the atom is still the same  element no matter how many neutrons it will just be an isotope.  

- Isotopes have essentially identical Chemical Properties but physical properties  (melting and boiling point) may differ slightly.

- The mass number (A) equals the number of protons Z and the number of  neutrons N.

- This is one way to distinguish isotopes  

- The number of neutrons can be determined from the known mass number and the  known atomic number.

Isotope Symbol Composition

- Mass number = protons and neutrons  

- Atomic number = protons

- Subscript = proton number Z

- Superscript = mass number A

- Because each element only has one atomic number sometimes is left out of the  symbol.

2.3 - Ions

- Ion - when an atom contains more or less electrons than protons and has a  charge.

Cation = a positively charged ion - contains less electrons than protons Anion = a negatively charged ion that contains more electrons than protons 2.4 Atomic Mass

- The AMU - atomic mass unit - scale uses carbon-12 as the standard to  express other masses of atoms

- Carbon-12 has 12 amu so one amu = 1/12 the mass of one carbon 12 atom  = 1.6606X10^-24 g

- The average mass of the individual isotopes is the Relative Atomic Mass 2.5 – The Periodic Table  

- On the periodic table, elements are in order of their atomic number form  smallest to largest  

- columns and rows emphasize periodic properties

- Horizontal row – a period – elements in the same period tent to have  properties that vary in a regular fashion – periods are labeled 1  through 7

- The stair step line separates metals on the left from the nonmetals on the right  

- A metalloid or semimetal is an element that has physical properties resembling a  metal but chemical reactivity more like a nonmetal – metalloids are along the  stair step line.

- Any element labeled with an A is a Main group element

- Any element labeled with a B is a transition metal

- The lanthanide and actinide series are the two lines on the bottom these elements  are the inner transition metals

- Alkali Metals = Group 1A (reactive)

- Alkaline Earth Metals = Group 2A (Less reactive)

- Halogens = Group 7A (diatomic molecules)

- Noble Gasses = Group 8A (single, inert atoms) - Noble gasses are inert, they do  not react chemically

- Diatomic Molecule – a molecule with two atoms (Have No Fear Of Ice Cold Beer)  – H, N, F, O, I, Cl, Br

- The position of an element on the periodic table helps to predict the charge on its  ion

- noble gasses do not form ions- they are the most stable -their stability is  associated with their number of electrons.

- Nonmetals usually gain electrons to form anions that have a noble gas electron  count

- Most main group elements lose electrons to form cations that have a noble gas  electron count

- Ions of the same group often form electrons of the same charge

Chapter 3  

3.1 – Ionic and Molecular Compounds

- For an electric current to go through a solution, ions must be present - Electrolyte – A substance that releases ions when it gets dissolved in water - Nonelectrolyte – A substance that does not release ions when dissolved in  water

Dissociation/Ionization – the process by which a compound dissolves in water to  produce ions.

- Strong electrolyte – dissociates easily and conducts electricity well. - Weak electrolyte – dissociates only partially and doesn’t conduct electricity  well.

- Nonelectrolyte - does not dissociate into ions at all and does not conduct  electricity.

Ionic Compound – Consists of oppositely charges cations and anions in proportions what are electrically neutral. Metal with a nonmetal usually. Molecular Compound – atoms of two or more nonmetals. – atoms do not  dissociate

- Much lower melting and boiling points than ionic compounds.

Binary compound – a compound containing atoms/ions of only two elements. Polar molecules – molecular compounds with very slight electrical charges – weak  – covalent bonds

- Acids do ionize in water even though they are molecular compounds – this is an exception.

- Acid’s dissociation in water is called ionization.

3.2 – Monatomic and Polyatomic Ions

- Monatomic Ion = an ion of a single atom

- The charges of ions for elements in the middle, especially transition metals,  cannot be predicted.

Nomenclature – they system of naming  

- Monatomic Anions – The first part of the element name (the root) and the  suffix -ide – ex) Sulfide ion = S^2-

- Monatomic Cations – are not given a suffix – they are just the name of the  element and the word ion ex) Sodium ion = Na^+ - Polyatomic Ions – an atom containing two or more atoms of usually more  than one element.

- The most common polyatomic ions have oxygen attached to them –  these are called Oxyanions – typically a combination of oxygen with  a nonmetal (can be a metal sometimes)

- Nearly all polyatomic ions are negatively charged – Exception =  NH^4+

- Rules for naming

- Greatest # of oxygen atoms – Prefix per + root name + suffix  ate

- 2nd Greatest # of Oxygen atoms – Root name + suffix ate  

- 2nd to Least # of Oxygen atoms – Root name + suffix ite

- Least # of Oxygen atoms – Prefix hypo + root name + suffix ite - If there are only 2 oxyanions only use the suffixes ate and ite, ate for  the greater number of oxygen atoms and ite for the lesser number. 3.3 – Formulas for Ionic Compounds

- Formula Unit – the smallest repeating unit of an ionic formula – has a net charge  of zero

- Chemical names start with the name of the element that is farther to the  left or farther down on the periodic table.

- The name of the metal is written first.

- The formula unit does not give the exact count of ions, it gives the  smallest ratio  

- Molecular formulas give the exact count of ions in the molecule. - Formulas for Ions containing polyatomic ions – sum of charges must be zero - Cations appear first in both the name and the formula.

- use parentheses around a compound that appears more than once to  balance a charge.

- Writing formulas from Ionic names  

1st determine if a polyatomic ion is present

2nd determine the charge on the ions that make up the compound – if the ion  is monatomic, look at the periodic table, if the ion is polyatomic,  we must recall its charge.

3rd write the correct formula

Ex) Calcium ion + Nitride Ion = Ca 2+ and N 3- = Ca3N2

3.4 - Naming Ionic Compounds  

- Binary Ionic Compounds – Cation then anion then -ide ending. – Prefixes are not  used.

- Naming Compounds containing Polyatomic Ions – Cation followed by the  name of the polyatomic ion.

- Many metals, especially transition metals, can exhibit multiple charges. - Stock System – using roman numerals  

- In parentheses after the metal will be a roman numeral to indicate  charge  

- Write the name of the metal, followed by the roman numeral, then  write the root of the nonmetal name with the -ide ending. - Writing formulas from names – Ex) iron(II) sulfide – iron has a charge  of plus three since it is the metal cation and sulfur has a  charge of 2- from the periodic table, so the formula is Fe2S3. - Older Method of Naming – the name of the metal is changes by adding  the suffix -ous or -ic to the root of its Latin name. - ous is for the lower of the two charges

- ic is foe the higher of the two charges

- This method only applies to metals that can exhibit 2 charges only. - Hydrates – ionic compounds that have stoichiometric amounts of water attached  to them and included in their formula units.

- Hydrates are names with prefixes (mono/di/tri etc.)

- Ex) MgSO4 7H2O = Magnesium sulfate heptahydrate  

3.5 – Naming and Writing formulas for Molecular Compounds  - Some nonmetals combine to form more than one compound 0 we must  identify the correct ratio of atoms in the molecule. - Naming – state the name of the element that is farther to the left (or  if two elements are in the same group use the one that is  farther down) on the periodic table.

- Add the suffix -ide to the end of the root name of the 2nd 

element.  

- Use Greek Prefixes to indicate the number of atoms in the  

compound.  

- The prefix mono is usually not used except in the cases  

of carbon monoxide nitrogen monoxide etc. - This is now the molecular formula that tells the actual number  of atoms present in a single molecule of a molecular  compound.

- Some binary molecular compounds are known by their trivial names,  or systematic names Ex) H2O = Water

3.6 – Acids and Bases  

- Acid = a substance that when dissolved in water provides hydrogen ions (H+).  This is called ionization when a compound that is not an ionic compound  provides ions when dissolved in water.

- Acids are molecular compounds but behave differently.

- The H+ ion does not exist in solution by itself it is usually surrounded  by water molecules that form a hydronium ion H3O+ - Organic acids are compounds that contain the combination of atoms C,  O2, and H – these are called Carboxylic acids. Ex) CH3CO2H - Not all compounds containing hydrogen are acids in water – only when  hydrogen combines with an element on the far right of the  periodic table (excluding noble gasses) the compound is an acid. - Writing Formulas for Acids – place the hydrogen first and write (aq) at the end  to indicate that it is an acid. Ex) HCl (aq)

- Hydrogens with a polyatomic ion can also be an acid. Ex) Nitric Acid HNO3 - Binary Acids in Solution – named with the prefix hydro- followed by the  stem of the nonmetal name with the suffix -ic and the word acid at  the end.  

- Naming acids with a polyatomic ion

 – The prefix hydro- is not used – remove the -ate ending from the  polyatomic ion and replace it with -ic.  - If the polyatomic ion has less oxygen atoms, replace the -ite ending  with -ous.  

-Base = a substance that reacts with an acid in an aqueous solution to form H2O. - Most common bases contain a Hydroxide ion (OH-) or can produce OH- ions  when in solution.

- Some Common Bases – NaOH, KOH, Mg(OH)2, NH3 ammonium. - Strong Acids and bases ionize/dissociate completely and are strong electrolytes. - Weak Acids and bases are wear electrolytes ex) Acetic Acid.

3.7 – Predicting Properties and Naming Compounds

- By deducing the elements in a compound from its name or formula, we can  predict many of its properties.  

- Ex) MgCl2 – it has a metal and a nonmetal so the compound is ionic which  must mean that it may be a brittle solid with a high melting and  boiling point and we can name it by ionic rules.

- Diprotic Acids = 2 hydrogen protons – Triprotic acids = 3 hydrogen  protons

- When writing a chemical formula, write the symbols in the order in which the  elements appear going from left to right on the periodic table. - If two elements are in the same column, write the name of the element that  is lower 1st.

- Hydrogen Exception – write H AFTER all elements EXCEPT those in groups 6A and 7A.  

Chapter 4  

4.1 – Percent composition – an expression of the proportion of the total mass  contributed by each element.

- Constant no matter the sample size – follows the law of definite proportions. - % element = (mass (g) element/ mass (g) sample) X 100%

- The percent masses of each element in the compound/ sample should all  add up to 100%.

- We can also find Mass from % composition. – Use 100g as a sample base - Mass of sample (g) X (% element changed into (g)/ 100g sample) =  Mass of element in the given sample.

4.2 – Mole Quantities  

- A Mole = 6.022 X 10^23 atoms/ions/molecules/formula units

- 6.022X 10^23 = Avogadro’s Number – the same as the number of atoms in  exactly 12g of carbon-12

- 1 Mole can contain different volumes and different masses of different  materials, but it is always the same amount of formula units  

(atoms/ions/molecules/formula units).

- To calculate the number of formula units, multiply the number of moles by  the conversion factor that cancels out moles using Avogadro’s number.   - 2 mole Cu X (6.022X10^23 formula units/ 1 mole Cu) =  

1.204X10^24 formula units Cu

- To calculate the number of atoms/ ions in 1 mole of an element, add a step  to the previous equation looking at the chemical formula to determine how  many atoms or ions are in 1 molecule or formula unit.

- 1 mole O2 X (6.022X10^23 molecules O2/ 1 mole O2) X ( 2 O atoms/  1 molecule O2) = 1.204 X 10^23 O atoms

- Molar Mass = (MM) The mass of 1 mole of a substance – Use the periodic  table to determine the molar mass of any element or compound. – If a  compound, the molar mass is equal to the sum of all relative atomic massed  of its component elements.

- Moles from Mass - Mass (g) of the substance X (1 mole Substance/ MM (g)  substance) = Moles Substance

- Mass from Moles – Moles substance X (MM (g) substance/ 1 mole substance) = Mass (g) substance

- Moles to Atoms – Moles element X (6.022X10^23 atoms element/1 mole  element) = Atoms element

- Molecules from Mass – 150g sample X (1 mole sample/ MM (g) sample) X  (6.022X10^23 molecules sample/ 1 mole Sample) = Molecules sample 4.3 – Determining Empirical and Molecular Formulas

- Empirical Formula (EF = my abbreviation) – expresses the simplest ratios of  atoms in a compound – written with the smallest whole number subscripts. - To find the EF, determine the smallest subscript and divide the other  subscripts by its value, or another least common factor.

- If the subscripts are not evenly divisible by a number other than 1,  then the EF and the molecular formula are not the same.

- The molecular formula is either the same as or a multiple of the EF. –  Molecular formulas contain more information than EFs.

- Determining Empirical formulas from data on the chemical composition of a  compound.  

- If given masses of elements in a compound, convert them to moles  using MM.

- We need to convert the number of moles into whole numbers  of atoms, so take the smallest number of moles of an element  

and divide the moles of the other elements by this number.  

- This will give you the number of moles of each element, which can be translated to subscript numbers in a chemical formula.

- Empirical Formulas from Percent Composition containing 2 elements and more  than 2 elements.

- Convert the percent composition of the elements in the compound to g of  that compound using 100g hypothetical sample.

- Convert the mass of each element to moles using MM.

- Divide each moles of element by the smallest of the mole quantities. - Now you have whole number moles that you can turn into subscripts for the  EF

- For more than 2 elements, just add in the step to convert the extra  elements.

- Empirical Formulas with fractional mole ratios.

- Decimal places may not always be close to whole numbers, thus rounding is not appropriate. Ratios may have fractional values like 1.25 or 1.5. In this  case, multiply each ratio by a small whole number to make all of the  subscripts whole numbers.

- Molecular Formulas from Empirical Formulas.

- We can calculate the mass of 1 mole of formula units using the EF. - If the experimental MM is the same as the calculated MM, the molecular  formula is the same as the EF.

- If the experimental MM is greater than the calculated MM of the formula, the molecular is some multiple of the empirical formula.

- Experimental MM/ EF MM (calculated) = The multiple that the  molecular formula is greater by.  

- Determining Percent Composition from Chemical Formula 

- If we know the chemical formula of a compounds, we can use the  relationship between moles and mass to determine its % composition without having to determine mass.

- To simplify converting a chemical formula into % composition, assume a  sample size of 1 mole.

- The mass of an element in 1 mole of the compound is equal to the number  of moles of the element multiplied by the MM of the element.

- The % composition of the element then equals the Mass of the element  divided by the MM of the whole compound, multiplied by 100% to get %  composition.

4.4 – Chemical Composition of Solutions

- Chemical reactions between solids are normally slow, so to speed up  reactions, we usually carry out such reactions by first dissolving the  compound in a liquid to form a solution.

- Solute = The substance being dissolved (present in a lesser amount) - Solvent = The substance doing the dissolving (present in a greater amount) - Concentration – The relative amounts of solute and solvent in the solution. - When comparing solutions, we can describe them as dilute or concentrated. - A dilute solution contains a relatively small amount of solute.

- A concentrated solution contains a relatively large amount of solute. - Indications of concentration that are not very accurate =  

Taste/color/density/viscosity by observation.

- Percent by Mass – the % by mass of a solution equals the mass of the solute divided by the mass of a solution multiplied by 100%.

- % mass = (Mass of solute/Mass of solution) X 100%

- The mass of the solution is the sum of the masses of solute and  solvent.

- This method does not depend on temperature. Volume may change  with temperature, but the mass does not. - % mass is also unaffected  by the presence of solute, unlike volume.

- Molarity (M) = The concentration of a solution can be expressed by molarity, the number of moles of solute dissolved in 1 L of solution.

- Molarity = Moles of solute/Liters of Solution

- Molarity from Mass  

- Find the moles of solute using MM.

- Convert volume to liters, if not already liters.

- Divide the moles of solute by the volume of the solution in liters to  get Molarity.

- Moles from Volume and Molarity 

- To get moles form volume, multiply by the molarity conversion factor  that cancels out units of liters and gives units of moles.

- Determine the moles of ions/atoms by looking at the chemical  formula and determine how many moles of that ion/atom are in the  formula unit.

- To get moles of the ion/atom, use the conversion factor that cancels  the appropriate unit. You are not left with moles of the specific  ion/atom in the formula unit/molecule.

- Mass from Volume and Molarity  

- If we know the volume and molarity of a solution, we can calculate  the mass of solute in that volume of solution

- To determine the mass of the solute, 1st determine the moles of solute by multiplying the molarity by the volume of solution in L to cancel out  L and be left with moles.

- Convert moles to mass using MM.

- Volume from Molarity and Mass  

- We can calculate the volume of a solution required to make a solution of a specific concentration from a given amount of solute.

- Given molarity and mass, convert mass to moles using MM. - Convert moles to volume of solution in L by multiplying the moles of  solute by the molarity conversion factor that cancels out moles and be  left with L.

- Dilution – to lower the concentration of a solution, we use the process of  dilution, adding more solvent to a solution. The relative numbers of solute  and solvent particles increase and the number of solvent particles increases  the volume, but the number of solute particles stays the same and now they  are spread out through a greater volume, so their concentration is less.

- If we dilute a solution of a known molarity and, we can calculate the  molarity of the diluted solution.

- Moles (concentrated) = Molarity (concentrated) X Volume  (concentrated)

- Moles (concentrated) = Moles (diluted)

- Moles (diluted) = Molarity (diluted) X Volume (diluted)

Therefore – M(dil) X V(dil) = M(conc) X V(conc)  

Rearrange this equation to solve for the Molarity of the diluted solution - M(dil) = (Mconc X Vconc)/V(dil)

- This formula can be rearranged to solve for any one of these  four variables.

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