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CHEM 1030 Dr. Streit- Week 1 Notes

by: Rachel Ferrell

CHEM 1030 Dr. Streit- Week 1 Notes CHEM 1030 - 003

Marketplace > Auburn University > Chemistry > CHEM 1030 - 003 > CHEM 1030 Dr Streit Week 1 Notes
Rachel Ferrell
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Covers chapter 1 and part of chapter 2
Fundamentals Chemistry I
John D Gorden
Class Notes
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This 5 page Class Notes was uploaded by Rachel Ferrell on Friday January 22, 2016. The Class Notes belongs to CHEM 1030 - 003 at Auburn University taught by John D Gorden in Fall 2015. Since its upload, it has received 181 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.


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Date Created: 01/22/16
Rachel  Ferrell   CHEM  1030-­‐003   1/19/16     Chapter  1:  The  Study  of  Chemistry   • Chemistry-­‐  the  study  of  matter  and  changes  matter  undergoes   • Matter-­‐  anything  that  has  mass  and  occupies  space   • Molecules-­‐  tiny  pieces  of  a  substance;  made  up  of  atoms   The  Scientific  Method:   • Scientific  method-­‐  a  set  of  guidelines  in  place  for  scientists  to  ensure  their  findings  can  be   widespread  and  valid   o 1)  Gather  data  via  observations  and  experiments   o 2)  Find  patterns  or  trends  in  collected  data   o 3)  Summarize  findings  with  a  law-­‐  a  concise  statement  that  makes  relations  between   phenomena   o 4)  Formulate  hypothesis-­‐  means  tentative  explanation   o 5)  Hypothesis  evolves  into  theory-­‐  unifying  principle  that  explains  a  large  set  of   experimental  observations;  can  also  predict  phenomena  that  have  not  happened  yet   Classification  of  Matter:   • classify  matter  as  either  a  substance  or  a  mixture   o Substance-­‐  a  form  of  matter  that  has  a  definite  composition  and  distinct  properties   § ex.  Salt  (NaCl),  water  (H2O),  carbon  dioxide  (CO2),  oxygen  (O2)   § these  elements  are  combined  chemically;  must  exist  together  to  keep  same   properties   § differ  in  substances  as  well  as  color,  taste,  smell,  etc.   o Mixture-­‐  a  physical  combination  of  2  or  more  substances   § Homogeneous  mixture-­‐  uniform  throughout;  solutions   • Ex.  Sea  water,  apple  juice   § Heterogeneous  mixture-­‐  not  uniform  throughout   • ex.  Trail  mix,  chicken  noodle  soup   • can  be  separated  by  physical  processes  that  do  not  change  identity  of   substance   • States  of  matter:   o Solids-­‐  particles  close  together  in  orderly  fashion;  does  not  conform  to  shape  of  container   o Liquids-­‐  particles  close  together  but  not  rigid  in  position;  does  conform  to  shape  of   container     o Gases-­‐  particles  far  apart;  conforms  to  shape  and  volume  of  container   o All  substances  can  exist  as  any  of  these  states  of  matter   o Changing  states  does  not  change  identity  of  the  substance   Properties  of  Matter:   • Quantitative  properties-­‐  measured/expressed  by  numbers   • Qualitative  properties-­‐  no  measurement;  based  on  observation   • Physical  properties:   o Can  be  observed  or  measured  without  changing  the  identity  of  the  substance   o Ex.  Color,  melting  point,  boiling  point,     o Physical  change-­‐  state  of  matter  changes,  but  the  identity  does  not  change;  like  changing   between  a  solid,  liquid  or  a  gas   • Chemical  properties   o A  property  shown  when  substance  reacts  with  another  substance;  it  is  no  longer  is  the   same  substance   o Ex.  Flammability,  corrosiveness   o Chemical  change-­‐  change  substance’s  composition,  original  substance  is  gone   o ex.  Digestion,  combustion,  oxidation   • Extensive  properties   o a  property  of  measured  value  that  depends  on  the  amount  of  matter   o ex.  Mass,  volume   • Intensive  Properties   o a  property  that  does  not  depend  on  the  amount  of  matter   o ex.  Temperature,  density   Scientific  Measurement:   • any  measured  quantity  must  have  a  unit   • SI  Base  Units   o Revised  metric  system  (International  system  of  units)   o Designed  to  be  universal  for  scientists   o 7  SI  Base  Units:   o Base  quantity   o Name  of  unit   o Symbol   o Length   o Meter   o m   o Mass   o Kilogram   o kg   o Time   o Second   o s   o Electric  current   o Ampere   o A   o Temperature   o Kelvin   o K   o Amount  of  substance o Mole   o mol   o Luminous  intensity   o Candela   o cd   o the  magnitude  of  a  unit  is  tailored  to  fit  a  particular  application  using  prefixes  to  the  base   units   o Tera-­‐   o T   o 1X10^12   o Giga-­‐                   o    1X10^9    G   o Mega-­‐   o M   o 1X10^6   o Kilo-­‐   o k   o 1X10^3   o Deci_   o d   o 1X10^-­‐1   o Centi-­‐   o c   o 1X10^-­‐2   o Milli-­‐   o m   o 1X10^-­‐3   o Micro-­‐   o μ   o 1X10^-­‐6   o Nano-­‐   o n   o 1X10^-­‐9   o Pico-­‐   o p   o 1X10^-­‐12   • Mass   o Measure  of  amount  of  matter  in  an  object  or  sample   o Weight-­‐depends  on  gravity   o Mass-­‐  does  not  change   o The  Atomic  Mass  Unit  (amu)=  used  to  express  the  mass  of  atoms  and  other  things  of   similar  size   § 1  amu=  1.6605378X10^-­‐24  g   § essentially-­‐  1  amu  is  very  small   • Temperature   o Celsius  Scale   § Freezing  point=  0°C   § Boiling  point=  100°C   o Kelvin  Scale   § The  “absolute  scale”   § Lowest  possible  temp=  0K  (absolute  zero)   § Kelvin  is  considered  an  SI  unit   o To  convert  between  C  and  K:   § K=  °C  +273.15   § C=    K-­‐273.15   o Fahrenheit  scale   § Freezing  point=  32°F   § Boiling  point=  212°F   § Good  for  body  temperature  because  1°F=9/5  °C;  therefore  making  it  more  of  an   exact  measurement   § Temp  in  °F=  (9/5  X  temp  in  °C)  +32°F     Rachel  Ferrell     CHEM  1030   1/21/16     Chapter  1  cont.     Derived  Units:  Volume  and  Density   • Volume/density-­‐  require  units  that  are  not  found  in  SI  units   • Must  use  SI  Units  in  combination  with  each  other   • SI  Unit  for  Volume=  m or  L  (liters  is  a  more  practical  measurement)   o 1  dm  =  1L   o 1  cm  =  1  mL       o the  units  for  volume  are  cubed  because  length  X  width  X  height of  a  cube   • Density   o Density  is  the  ratio  of  mass  to  volume   ▯ o d  =       Also  note  that  v▯  and  m=d  x  v   ???? o d=  density   o m=mass   o v=volume   3 o SI  Derived  unit=  kg/m   o Common  units  are  based  on  states  of  matter:   § g/cm =  solids   § g/mL  =  liquids   § g/L  =  gases   Uncertainty  of  Measurement:   • 1)  Exact  Measurement:   o =those  that  have  defined  values   § 1  kg  =  1000g   § 1  dozen  =  12  objects   o =those  determined  by  counting   § 28  students  in  the  class   • 2)  Inexact  Measurement:   o measured  by  anything  other  than  counting;  accounts  for  human  and  machine/tool  error   § ex.  Length,  mass,  volume,  time,  speed,  etc.   § anything  you  have  to  use  a  tool  to  measure   o an  inexact  number  must  be  reported  to  show  uncertainty         Significant  Figures  (aka  Sig  Figs):   • =meaningful  digits  in  a  set  of  reported  numbers   o the  last  digit=  uncertain  digit   • ***Sig  fig  Rules***   o 1)  Any  nonzero  digit  is  significant   § 112.1=  4  sig  figs   o 2)  Zeros  between  nonzero  digits  are  significant   § 305  =  3  sig  figs   § 50.08  =  4  sig  figs   o 3)  Zeros  to  the  left  of  the  first  nonzero  digit  are  not  significant   § 0.0023  =  2  sig  figs   § 0.00000001  =  1  sig  fig   o 4)  Zeros  to  the  right  of  the  last  nonzero  digit  are  significant  if  decimal  is  present   § 1.200  =  4  sig  figs   o 5)  Zeros  to  the  right  of  the  last  nonzero  digit  in  a  number  with  no  decimal  may  or  may  not   be  significant   § 100  =  1,  2,  or  3  sig  figs   § to  avoid  this  ambiguity,  write  these  types  of  numbers  in  scientific  notation,  so  that   no  extra  zeros  are  present     • Sig  Figs:  Adding  and  Subtracting:   o Answer   can’t   have   more   digits   to   the   right   of   decimal   point   than   any   of   the   original   numbers   o      102.50     +        0.231          102.731  Round  to  102.73*       o   • Sig  Figs:  Multiplication  and  Division   o Determined  by  original  number  with  the  lowest  number  of  sig  figs   o 1.4  X  8.011  =  11.2154→round  to  11  (2  sig  figs)   o 11.57  /  305.88  =  0.0378252→  round  to  0.03783  (4  sig  figs)   • Exact  numbers  can  be  considered  to  have  infinite  number  of  sig  figs   o If  there  are  exact  numbers  in  a  problem,  do  not  use  them  do  determine  number  of  sig  figs   • Sig  figs  only  apply  to  inexact  numbers  (numbers  that  are  not  counted)     Rounding  Rules:   •  Round  only  after  all  steps  of  the  problem  have  been  completed   • <5→round  down   • ≥  5  →  round  up   Accuracy  and  Precision:   • Accuracy-­‐  tells  us  how  close  a  measurement  is  to  the  true  value   • Precision-­‐  tells  us  how  close  a  series  of  replicate  measurements  are  to  one  another   o ex.    1)  you  hit  3  darts  on  the  bull’s-­‐eye→  good  accuracy  and  precision   o              2)  You  hit  3  darts  to  the  right  of  the  bull’s-­‐eye→bad  accuracy,  good  precision   o              3)  You  hit  3  darts  in  3  different  places  not  on  the  bull’s-­‐eye→  bad  accuracy,  bad   precision         Using  Units  and  Solving  Problems:   • conversion  factor=  a  fraction  in  which  same  quantity  is  expressed  one  way  in  the  numerator  and   another  in  the  denominator   o ex.       ▯  ▯▯    OR         ▯.▯▯  ▯▯ ▯  ▯▯ • dimensional  analysis=  use  of  conversion  factors  in  a  problem;  factor  label  method                  


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